DESCiMMi.c  ilNORGANIC 


GENERAL  CHEMISTRY 


Eext^lSooft  for  Colleges 


BY 


PAUL    C.    FREER,    M.D.,    PH.D.    (Munich} 

PROFESSOR   OF   GENERAL    CHEMISTRY,   AND    DIRECTOR   OF   THE   LABORATORY    OF 
GENERAL    CHEMISTRY,    UNIVERSITY    OF    MICHIGAN 


REVISED  EDITION 


i  ITS  asms? 


Boston 

ALLYN     AND     BACON 

1895 


COPYRIGHT,  1894, 

BY 
ALLYN    AND    BACON. 


I 


C.  J.  FETEBS  &  SON, 

TYPE-SETTEB6  AND  ELECTBOTYPEBS, 

145  HIGH  STBEET,  BOSTON. 


PEE  FACE. 


IN  compiling  the  descriptive  portions  of  this  work,  I  have  chiefly 
used  GRAHAM-OTTO'S  "  Lehrbnch  der  Allgemeinen  Chemie  "  (last 
edition)  and  LADEXBURG'S  "  Handworterbuch  der  Chemie,"  although, 
wherever  any  of  the  facts  which  it  was  necessary  to  incorporate 
seemed  doubtful,  the  original  sources  in  the  chemical  periodicals 
have  been  consulted ;  of  course,  the  discoveries  which  have  been 
brought  forth  in  the  last  few  years  are  taken  entirely  from  the 
latter.  In  the  discussion  of  the  double  halides,  of  fluosilicic  acid, 
and  of  similarly  constituted  bodies,  I  have  adopted  the  views  which 
have  recently  been  brought  into  prominence  by  the  publications  of 
Prof.  Ira  Eemsen,  both  in  his  larger  text-book  and  in  his  contribu- 
tions to  current  chemical  literature.  In  those  portions  of  the  work 
which  refer  to  the  application  of  physical  methods  in  the  study  of 
chemistry,  OSTWALD'S  "  Outlines  of  General  Chemistry  "  and  LOTHAR 
MEYER'S  "  Theoretische  Chemie  "  have  been  consulted  and  not  infre- 
quently quoted.  My  views  upon  the  subject  of  valence  and  the 
use  of  structural  formulae  may  possibly  be  regarded  by  many  of 
my  colleagues  as  too  conservative;  but  I  have  been  led  to  adopt 
these  views  by  the  growing  conviction  that  the  dogmatic  use  of 
supposed  laws  of  valence  and  of  constitutional  formulse  founded 
upon  very  incomplete  experimental  evidence,  is  causing  more  harm 
than  good  to  the  advancement  of  chemical  science.  In  discussing 
chemical  changes  and  reactions,  I  have  endeavored  to  present  the 
various  topics,  not  as  a  series  of  isolated  facts,  but  as  so  connected, 
the  one  with  the  other,  that  there  is  scarcely  any  one  of  the  numer- 
ous phenomena  which  are  mentioned  in  this  work  which  does  not 
find  its  analogon  in  some  other  portion  of  the  field  of  chemical 
study.  The  attempt  has  been  made  especially  to  call  attention  to 
the  influence  exerted  by  the  nature  of  the  elements  which  make  up 
a  chemical  compound  upon  the  character  of  that  compound  itself. 
Tracing  those  connections  may  possibly  have  led  me  somewhat  into 

iii 


IV  PREFACE. 

the  realm  of  speculation,  notably  so,  perhaps,  in  my  endeavor  to 
explain  the  behavior  of  the  hydrogen  compounds  of  the  not-metals 
by  taking  into  consideration  the  relative  influence  exerted  by  the 
masses  of  the  atoms  which  go  to  make  the  molecules.  I  hope, 
however,  that  the  new  arguments  ventured  on  during  the  progress 
of  this  work  will  not  be  condemned  without  a  hearing.  Of  course, 
a  very  complete  knowledge  of  descriptive  chemistry,  both  inorganic 
and  organic,  is  necessary  before  the  study  of  so-called  physical 
chemistry  can  be  pursued  with  profit ;  nevertheless,  wherever  it 
has  seemed  to  me  that  some  elementary  facts  from  the  realm  of 
physical  chemistry  would  be  comprehended  by  the  pupil  taking  up 
beginning  chemistry,  I  have  not  hesitated  to  introduce  the  latter,  at 
the  same  time  giving  references  to  the  best  of  the  smaller  text- 
books on  the  subject.  The  atomic  weights  which  I  have  used  are 
taken  from  the  table  recently  prepared  by  F.  W.  Clarke,  with  the 
atomic  weight  of  oxygen  "  en  as  the  standard)  placed  at  16. 

The  laboratory  notes  j  appendix  cover  only  the  ground 

taken  by  the  not-metals ;  tney  are  not  intended  as  a  laboratory  man- 
ual, but  mainly  as  a  guide  to  both  teacher  and  pupil  in  compiling  a 
list  of  experiments.  Every  teacher  prefers  using  his  own  methods 
for  laboratory  instruction,  with,  of  course,  his  own  selection  of  the 
work  to  be  pursued ;  in  my  own  laboratory  I  follow  a  manual  which 
is  made  up  of  brief  directions,  accompanied  by  a  very  complete  set 
of  questions,  and  all  of  the  latter  must  be  answered  by  the  pupils. 
I  do  not  think,  however,  that  pupils  should  be  left  in  the  laboratory 
without  other  than  a  printed  guide ;  far  from  it,  I  like  to  see  the 
instructor  always  present  in  the  room  during  laboratory  hours, 
guiding  and  assisting  his  pupils,  and  not  infrequently  working 
with  them. 

Probably  this  work  cannot  be  advantageously  employed  in  the 
secondary  schools  ;  indeed,  it  is  adapted  for  the  use  of  students  who 
already  have  some  knowledge  of  the  elementary  principles  of  the 
science.  Beginners  should  be  taken  through  a  course  in  which  only 
a  few  elements  and  compounds  are  discussed,  with  the  purpose  of 
familiarizing  the  pupils  with  the  fundamental  laws  which  govern 
chemical  change.  During  the  progress  of  such  work  as  this,  I 
would  not  advise  the  use  of  chemical  symbols  or  any  reference  to 
the  atomic  theory.  Our  chemical  symbols  and  equations  are  in 
existence,  in  their  present  state,  only  because  of  the  difficult  experi- 
mental work  which  has  finally  succeeded  in  establishing  a  consis- 


PREFACE.  V 

tent  table  of  atomic  weights.  It  is  manifestly  impossible  to  make  a 
student,  without  experimental  knowledge,  understand,  in  all  its 
bearings,  a  theory  which  it  has  taken  some  ninety  years  to  place 
upon  its  present  footing.  If  an  elementary  course,  in  which  every 
stated  fact  has  been  proved  by  actual  experiment,  precedes  the  work 
given  in  this  book,  the  pupil  will  then  be  amply  fitted  to  look  at 
chemical  phenomena  from  the  basis  of  the  atomic  theory.  It  is  in 
the  hope  that  such  a  preparatory  course  has  gone  before,  that  I  have 
begun  this  text-book  with  the  atomic  theory. 

PAUL  C.  FREER. 
ANN  ARBOR,  June,  1894. 


After  the  first  five  hundred  copies  were  printed,  the  book  was 
subjected  to  a  careful  reading  by  several  chemists,  so  that,  it  is 
hoped,  all  misprints  and  errors  have  .<i  entirely  eliminated. 

PAUL   C.  FREEH. 
ANN  ARBOR,  May,  1895.  j 


\ 


CHAPTER  PAGE 

I.     Introductory 1 

II.     Oxygen 18 

III.     Hydrogen 27 

IY.     Water 36 

V.     Ozone  and  Hydrogen  Dioxide 47 

VI.     The  Halogen^ 53 

VII.     Fluorine  and  Hydrofluoric  Acid 55 

VIII.    jChlorjjie 58 

IX.     Hydrochloric  Acid 66 

X.     Bromine  and  Hydrobromic  Acid 78 

XL     Iodine  and  Hydroiodic  Acid 83 

XII.     fhTOxygen  Family 89 

XIII.  Sulphur 91 

XIV.  Hydrogen  Sulphide 96 

XV.     Selenium  and  Hydrogen  Selenide 102 

XVI.     Tellurium  and  Hydrogen  Telluride;  Comparative  Table  of 

the  Elements  of  the  Oxygen  Family 104 

XVII.     Valence  and  the  Oxygen  Compounds 107 

XVIII.     The  Compounds  of  Chlorine  with  Oxygen,  and  with  Oxy- 
gen and  Hydrogen 119 

XIX.     Compounds  of  Bromine  and  of  Iodine  with  Oxygen  and 

Hydrogen,  the  Compound  of  Iodine  with  Oxygen,  and 

the  Compounds  of  the  Halogens  with  each  other      .     .  129 

XX.     The  Compounds  of  the  Elements  of  the  Sulphur  Family  with 

Oxygen,  and  with  Oxygen  and   Hydrogen.      Sulphur 

Dioxide  and  Sulphurous  Acid 134 

XXI.     Sulphur  Trioxide,  Sulphuric  Acid,  and  the  remaining  Sul- 
phur Acids 144 

XXII.     The  Compounds  of  Selenium  and  Tellurium  with  Oxygen, 

and  with  Oxygen  and  Hydrogen .160 

XXIII.  Nitrogen  and  the  Atmosphere 163 

XXIV.  Compounds  of  the  Elements  of  the  Nitrogen  Family  .     .     .  175 

vii 


Vlll 


CONTENTS. 


CHAPTER  PAGE 

XXV.     Ammonia  and  the  other  Compounds  of  Nitrogen  and  Hy- 
drogen      182 

XXVI.     The  Compounds  of  Nitrogen  with  Oxygen,  and  with  Oxygen 

and  Hydrogen 195  ' 

XXVII.     Phosphorus  and  Phosphine 211 

XXVIII.     The  Compounds  of  Phosphorus  with  the  Halogens,  and  with 

Oxygen  and  the  Halogens 219 

XXIX.     The   Compounds   of  Phosphorus  with  Oxygen,  and   with 

Oxygen  and  Hydrogen 223 

XXX.     Arsenic  and  Arsine 234 

XXXI.     The  Compounds  of  Arsenic  with  the  Halogens,  with  Oxygen, 

and  with  Oxygen  and  Hydrogen 238 

XXXII.     The  Compounds  of  Arsenic  with  Sulphur,  and  with  Sulphur 

and  Hydrogen 244 

XXXIII.  Antimony  and  Stibine.     The  Compounds  of  Antimony  with 

the  Halogens | 248 

XXXIV.  The  Compounds  of  Antimony  with  Oxygen,  and  with  Oxygen 

and  Hydrogen.     The  Sulphides  of  Antimdny  ....  254 
XXXV.     Bismuth.     The  Compounds  of  Bismuth  with  the  Halogens, 
with  Oxygen,  with  Oxygen  and  Hydrogen,  and  with 

Sulphur 257 

XXXVI.     The  Elements  of  the  Carbon  Family 265 

XXXVII.     Carbon 269 

XXXVIII.     The  Compounds  of  Carbon  with  Hydrogen 274 

XXXIX.     The  Compounds  of  Carbon  with  Chlorine,  with  Chlorine 

and  Oxygen,  with  Oxygen,  and  with  Sulphur ....  285 
XL.     Compounds  of  Carbon  with  Nitrogen,  with  Nitrogen  and 

Hydrogen,  and  with  Nitrogen,  Oxygen,  and  Hydrogen,  294 
XLI.     Silicon,  the   Compounds  of   Silicon   with   Hydrogen,   and 

with  the  Halogens,  the  Oxide,  and  Acids  of  Silicon    .  300 

XLII.     Germanium  and  its  Compounds 309 

XLIII.     Tin  and  its  Compounds 312 

XLIV.     Lead  and  its  Compounds 320 

XLV.     The  Elements  of  the  Boron  Family  (The  Earths)  ....  326 

XL VI.     Boron  and  its  Compounds 328 

XLVII.     Aluminium  and  its  Compounds 333 

>   XLVIII.     Gallium,  Indium,  and  Thallium 343 

^  XLIX.     The  Determination  of  Atomic  Weights.    Dulong  and  Petit' s 

Law.     The  Law  of  Isomorphism 347 

L.     The  Periodic  System  of  the  Elements 361 


CONTENTS. 


IX 


XLIH. 

\  LIV. 

LY. 

^ 


CHAPTER  PAGE 

^*"  LI.     Neutralization.      Double  Decomposition.     Dissociation   of 

Electrolytes  ................  375 

The  Alkali  Metals      ..............  383 

Copper,  Silver,  and  Gold     ............  396 

The  Family  of  the  Alkaline  Earths  .........  411 

Zinc,  Cadmium,  and  Mercury  ......  •    .....  423 

The  Elements  belonging  to  the  Primary  Groups  of  the 

Families  III.,  IV.,  and  V.,  of  the  Long  Periods  .  .  .  437 
The  Elements  belonging  to  the  Primary  Group  of  the  VI. 

Family  .................  443 

The  Element  forming  the  Primary  Group  of  the  VII. 

Family   .................  459 

Iron,  Cobalt,  and  Nickel     ............  470 

The  Remaining  Elements  of  the  VIII.  Family.  (The  Plati- 

num Group.  )  ...............  489 

APPENDIX  OF  LABORATORY  NOTES  ......  497 


LVII. 
LVIII. 

LIX. 
LX. 


GENERAL    CHEMISTRY. 


CHAPTEE    I. 

INTRODUCTORY. 

THE    ATOMIC    THEORY    AND    THE    COMPOSITION    OF    CHEMICAL 
COMPOUNDS. 

AMONG  the  most  important  theories  of  modern  physical  science 
is  the  one  which  is  based  upon  the  supposition  that  all  substances 
are  made  up  of  small  particles  called  atoms.  The  theory  that 
matter  is  not  infinitely  divisible,  but  that,  upon  attempted  separa- 
tion into  smaller  parts,  a  mass  not  capable  of  further  subdivision 
would  result,  was  held  by  the  Greek  and  Koman  philosophers  — 
by  Democritus,  Aristotle,  Epicurus,  and  Lucretius  —  and  has  been 
transmitted  to  the  present  generation  with  many  important  modifi- 
cations. The  idea,  for  we  could  scarcely  dignify  by  the  name  of 
theory  that  which  had  so  little  foundation  in  experiment,  was 
partially  lost  sight  of  during  the  dark  ages  of  chemistry,  —  during 
the  time  of  the  alchemists,  when  the  sole  aim  of  chemical  study 
was  mercenary,  when  scepticism  on  the  one  hand  and  popular 
superstition  on  the  other  had  stifled  all  originality  of  thought  and 
co-ordination  of  theory  in  this  field  of  knowledge ;  it  suffered  no 
better  fate  at  the  hands  of  those  who  succeeded  the  alchemists,  for 
they  were  men  who  used  their  small  knowledge  of  chemical  facts 
for  the  purpose  of  discovering  new  drugs  and  remedies ;  it  could 
expand  into  what  it  is  only  when  chemistry,  freed  from  the  bane 
of  superstition,  began  to  be  followed  for  the  sole  purpose  of  increas- 
ing human  knowledge. 

We  trace  the  growth  of  a  science  of  chemistry  from  the  begin- 
ning of  the  eighteenth  century,  for  then  chemists  began  to  have 
theories  founded  on  experiment;  these  Avere  undoubtedly  often 
false  and  misleading,  but,  nevertheless,  scientific  progress  was  inevi- 

1 


2  ATOMIC    THEORY;    HISTORY. 

table,  because  of  the  attempts  to  answer  the  problems  arising.  In 
this  century  fell  many  of  the  greatest  discoveries  of  modern  chem- 
ical science  /notably  the  proof  of  the  existence  of  more  than  one 
variety  of  gas  and  that  of  the  formation  of  the  atmosphere  from 
two  kinds  of  matter,  oxygen  and  nitrogen.  *It  was  learned  that  water 
could  be  produced  by  the  union  of  oxygen  and  hydrogen,  and  that 
substances,  in  burning,  absorb  a  constituent  of  the  atmosphere,  while 
in  so  doing  they  gain  in  weight,  the  gain  in  weight  of  the  burning 
substance  being  exactly  equal  to  the  loss  in  weight  sustained  by 
the  atmosphere^  To  the  knowledge  won  at  this  time  we  owe  our 
understanding  of  a  principle  of  nature  upon  which  all  chemical 
speculations  are  based;  that  of  the  conservation  of  matter.  The 
English  chemists,  Black,  Priestley,  and  Cavendish,  were  the  men 
whose  efforts  developed  so  many  new  facts;  but  it  was  owing  to 
the  clear  insight  into  the  meaning  of  these  discoveries  obtained  by 
Antoine  Laurent  Lavoisier,  that  a  greater  service  was  rendered  to 
humanity,  because  by  him  a  proper  explanation  of  the  phenomena 
involved  was  given.  Without  the  discovery  of  oxy_gen  by  Priestley, 
or  the  composition  of  water  by  Cavendish,  Lavoisier  might  not 
have  proved  the  law  of  the  conser^ajioji_oj_jnatter  nor  have  estab- 
lished the  theory  of  combustion  held  at  the  present  time ;  but  it  is 
equally  true  that,  without  Lavoisier's  genius  the  work  of  the  Eng- 
lish scientists  would  not  have  accomplished  the  result  of  preparing 
chemistry  for  the  unprecedented  advance  recorded  of  it  in  the  nine- 
teenth century.  During  the  time  of  these  great  discoveries  the 
atomic  jheory,  though  tacitly  accepted,  was  not  made  the  basis  of 
investigation ;  but  when  the  present  century  dawned,  chemists 
began  to  feel  the  need  of  some  rational  explanation  of  those 
phenomena  which  most  concerned  them. 

The  first  decade  brought  the  discovery  that  when  two  substances 
unite  chemically,  a  compound  is  always  formed  in  unvarying  pro- 
portions by  weight.  Thus,  iron  and  sulphur  unite  to  form  iron 
sulphide,  a  substance  in  which  there  are  four  parts  by  weight  of 
sulphur  for  every  seven  parts  by  weight  of  iron,  no  matter  where 
or  how  the  combination  takes  place  ;  and  there  is  also  another  com- 
pound of  iron  and  sulphur  in  which  seven  parts  by  weight  of  iron 
accompany  eight  of  sulphur.  Under  whatever  conditions,  or  in 
whatever  place,  either  of  these  chemical  bodies  are  produced,  the 
resulting  proportions  are  always  the  same.  If  there  is  more  sulphur 


LAW    OF   DEFINITE   PROPORTIONS.  3 

present  than  is  necessary  for  combination,  then  the  excess  of  sul- 
phur remains  unchanged,  and  if  more  iron  is  employed,  then  iron  is 
found  after  the  union.  We  have  two  compounds  of  carbon  and 
oxygen,  called  oxides  for  reasons  similar  to  those  which  gave  the 
name  of  sulphides  to  the  compounds  of  iron  and  sulphur.  In  one 
of  these,  six  parts  by  weight  of  carbon  are  united  with  eight  parts 
of  oxygen ;  in  the  other,  six  parts  of  carbon  are  united  with  sixteen 
parts  of  oxygen.  What  is  true  as  regards  iron,  carbon,  oxygen,  and 
sulphur  characterizes  the  multitude  of  other  substances  which  have 
been  studied  with  the  object  of  ascertaining  the  relative  proportions 
by  weight  in  which  the  constituent  parts  unite.  To  the  discoveries 
outlined  above,  have  been  added  the  demonstration  that  the  relative 
proportions  of  sulphur  and  oxygen,  for  instance,  are  preserved  in 
whatever  compound  they  are  encountered.  Thus,  sulphur  and 
oxygen  form  two  compounds,  called  oxides  of  sulphur;  in  one  of 
these,  two  parts  by  wreight  of  sulphur  unite  with  two  of  oxygen, 
in  the  other  two  of  sulphur  with  three  of  oxygen  ;  but  the  relation- 
ship in  the  weights  of  oxygen  and  sulphur  in  the  various  compounds 
cited  first  becomes  apparent  if  we  calculate  the  weights,  placing  sul- 
phur at  16,  while  preserving  the  proportion  between  the  various 
parts,  thus :  — 

Iron  and  sulphur 28  parts  of    iron  unite  with  16  of  sulphur. 

a  44  it  no          ..  ..         ..  ..  . .        oo     . .  . . 

Carbon  and  oxygen  ......       6      "       "  carbon    "        "       8  "    oxygen. 

It  it  it  a  4.  ..  ..  ..  ..  "JiJ       44  44 

Sulphur  and  oxygen      ....     16      "       "  sulphur "        "     16  "         " 
" 16      "       "       "         "        "      24  " 

These  relations  represent  actual  facts,  whatever  explanation  we 
may  see  fit  to  attach  to  them  ;  but  such  facts  as  these  necessarily 
give  rise  to  speculations  as  to  the  underlying  causes.  Why  should 
it  not  be  possible  to  have  28  parts  of  iron  united,  at  one  time  with 
16  parts  of  sulphur,  at  another  with  17  parts,  at  yet  another  with 
15  parts  ?  As  chemists  could  see  no  reason  for  such  regularity  in 
the  composition  of  matter,  the  facts  themselves  were  at  first  dis- 
puted, until  repeated  experiment  rendered  them  incontrovertible. 
Assuming  the  constancy  of  proportion  in  chemical  compounds,  even 
before  such  constancy  was  proved,  the  English  chemist,  John 
Dalton,  sought  an  explanation  in  the  following  hypothesis,  which 
has  been  accepted  as  a  basis  for  chemical  speculation  ever  since  its 


4  MODERN  ATOMIC  THEORY. 

establishment,   and  which  is  here  given  in  the  form  at  present 
accepted. 

Matter  is  not  infinitely  divisible,  but  is  composed  of  very  small 
and  discrete  entities  called  atoms,  there  being  as  many  different 
kinds  of  atoms  as  there  are  varieties  of  substance  which  have  never 
been  decomposed  into  two  or  more  forms  with  differing  properties. 
The  elements  having  weight,  the  atoms,  being  portions  thereof,  ne- 
cessarily also  have  weight ;  and  we  assume  the  weight  of  an  atom 
of  a  given  element  to  be  equal  to  that  of  each  other  atom  of  the 
same  element,  but  to  differ  from  that  of  an  atom  of  any  other 
element.  The  atoms  of  different  elements  unite  to  form  the 
smallest  individual  group  of  a  compound,  and  atoms  of  the  same 
kind  unite  to  produce  the  smallest  individual  group  of  an  element. 
These  groups  are  known  as  molecules,  and  the  agglomeration  of 
molecules  forms  tangible  matter.  The  weight  of  any  given 
molecule,  called  its  molecular  weight,  is  therefore  equal  to  the  sum 
of  the  weights  of  its  constituent  atoms.  If  I  subdivide  any  com- 
pound body,  water  for  instance,  I  can  continue  the  operation  until 
I  arrive  at  the  smallest  individual  particle  thereof,  a  molecule ;  if  I 
divide  this,  I  no  longer  have  water,  but  two  different  kinds  of  mat- 
ter, hydrogen  and  oxygen.  This  illustration  will  also  serve  to  show 
the  difference  between  a  so-called  chemical  and  a  physical  change. 
Water  can  be  decomposed  into  its  molecules  with  comparative  ease ; 
by  changing  it  into  steam  these  particles  are  so  far  separated  that 
they  travel  in  right  lines  independently  of  each  other.  A  much 
greater  heat  than  is  necessary  for  the  production  of  steam  (or  the 
application  of  some  other  form  of  energy,  such  as  electricity)  is 
necessary  to  effect  any  further  change,  and  this  change  brings  with 
it  the  destruction  of  the  nature  of  the  substance  in  question.  Water 
is  no  longer  present ;  but  in  its  place  we  have  two  different  kinds  of 
matter,  hydrogen  and  oxygen,  so  that  a  chemical  change  has  been 
brought  about. 

With  the  atomic  hypothesis  in  view,  the  constant  composition 
by  weight  of  compound  substances  is  readily  explained.  For 
instance,  a  molecule  of  iron  sulphide  is  composed  of  atoms  of  iron 
and  of  sulphur,  each  molecule  containing  the  same  number  of  atoms 
of  the  two  elements.  If  all  atoms  of  iron  are  alike  in  weight,  and 
if  all  atoms  of  sulphur  bear  the  same  relationship  to  each  other,  it 
follows  that  every  molecule  of  iron  sulphide  must  have  the  same 


LAW   OF   MULTIPLE  PROPORTIONS.  5 

composition  as  every  other  molecule  of  the  same  substance;  and 
from  this  it  follows  that  tangible  quantities  of  iron  sulphide,  which 
are  simply  agglomerations  of  the  individual  molecules,  must  have 
the  same  proportional  composition  by  weight  as  these  individual 
particles.  By  accepting  the  theory  as  it  has  been  outlined,  the 
unvarying  composition  of  purely  chemical  compounds  necessarily 
follows.  Of  course,  two  or  more  substances  may  be  mixed  in  any 
proportion,  but  such  a  mixture  does  not  have  the,  characteristics 
of  a  chemical  compound.  The  various  constituents  of  a  mixture 
can  be  separated  with  greater  or  less  ease  by  simple  mechanical 
operations ;  but  as  soon  as  a  chemical  compound  is  formed  from  the 
various  parts  of  any  mixture,  a  substance  having  a  definitely  con- 
structed molecule  results. 

We  saw  that  iron  and  sulphur  are,  however,  capable  of  forming 
two  compounds  with  each  other.  In  comparing  these  two  sulphides, 
it  has  been  discovered  that  the  part  by  weight  of  sulphur  which 
unites  with  a  given  weight  of  iron  in  one  of  these  substances  is 
exactly  twice  that  which  is  found  united  to  the  same  quantity  of 
iron  in  the  other ;  and  repeated,  painstaking  experiments  have  dis- 
covered a  great  number  of  similar  cases,  in  which  one  element  forms 
tAvo  or  more  distinct  compounds  with  some  other.  In  comparing 
such  compounds  it  has  ahvays  been  shown  that  the  parts  by  weight 
of  one  of  the  elements,  Avhen  united  to  a  fixed  quantity  by  Aveight 
of  the  other,  bear  a  simple  relationship  to  each  other;  for 
example :  — 

Mercury  and  oxygen  form  two  oxides ;  in  one,  for  every  100  parts  of  mer- 
cury there  are  4  parts  of  oxygen ;  in  the  other,  for  every  100  parts  of  mercury 
there  are  8  parts  of  oxygen. 

Nitrogen  and  oxygen  form  five  oxides ;  in  these  compounds  the  parts  by 
weight  of  oxygen  which  are  united  with  100  parts  by  weight  of  nitrogen 
are  as  57.  1  :  114.2  :  171.3  :  228.4  :  285.5,  which  figures  are  to  each  other  as 
1  :2  :3  :"4  :  5. 

The  results  of  these  discoveries  can  be  summed  up  in  the 
folloAving  law  of  .multiple  proportions :  — 

If  two  elements,  a  and  b,  unite  in  more  than  one  proportion,  the 
parts  by  weight  of  b  Avhich  Avill  unite  Avith  a  definite  quantity  of  a 
Avill  be  in  simple  ratio  to  each  other.* 

*  See  John  Dalton,  New  System  of  Chemical  Philosophy  (1808). 


6  STOICHIOMETEIC   QUANTITIES. 

This  law  of  multiple  proportions  is  readily  explained  by  the 
atomic  hypothesis.  For,  let  us  suppose,  using  the  two  sulphides  of 
iron  as  an  example,  that  the  one  composed  of  28  parts  of  iron  to  16 
parts  of  sulphur  has  a  molecule  constructed  of  one  atom  of  iron 
and  one  atom  of  sulphur.  In  order  to  change  this  molecule  into  one 
containing  more  sulphur,  the  only  possible  means  is  by  the  addition 
of  another  atom  of  sulphur.  But  as  the  atoms  of  sulphur  all  have 
the  same  weight,  it  follows  that  the  amount  of  sulphur  in  the 
newly  formed  molecule  must  be  to  that  in  the  original  as  2  :  1.  We 
might  represent  the  change  graphically  as  follows,  using  the  black 
circle  to  represent  an  atom  of  iron,  the  white  one  sulphur. 

28  16  16  16  28  16 

•   O  +  O  000 

Molecule  of  iron  sulphide  +  1  atom  of  sulphur  =  Molecule  No.  2.  of  iron 
sulphide. 

We  could  have  come  to  the  same  conclusions,  deducing  the  law 
of  multiple  proportions  as  a  necessary  consequence  of  the  atomic 
structure  of  matter,  had  we  used  combinations  of  any  other 
elements ;  for  the  law  is  universal  in  its  application.  What  is  true 
of  the  individual  molecule  must  also  be  true  of  tangible  matter. 

The  proportions  by  weight  in  which  the  elements  unite  are 
called  their  sto'ichiometric  quantities  ;  these  belong  not  only  to 
individual  pairs  of  elements,  but  are  universal.  This  truth  at 
once  becomes  apparent  if  we  select  a  given  weight  of  some  one 
element  and  then  determine,  experimentally,  the  parts  by  weight 
in  which  all  other  elements  will  unite  with  it :  by  this  means  it 
will  be  found  that  the  quantities  in  which  the  various  elements  unite 
with  the  standard  are  also  the  relative  proportions,  or  multiples  or 
submultiples  of  the  proportions,  in  which  they  will  unite  ivith  each  other. 
These  facts  are  in  entire  accordance  with  the  atomic  theory  j  for  as 
all  chemical  reactions  take  place  between  the  atoms,  and  as  the 
atoms  of  a  given  element  have  a  constant,  equal  mass,  different  from 
the  masses  of  the  atoms  of  all  other  elements,  it  follows  that  the 
relative  proportions  between  the  parts  by  weight  in  which  the 
elements  combine  must  also  be  constant.  Were  all  chemical  com- 
pounds to  be  of  the  simplest  nature,  so  that  each  molecule  would  be 
formed  by  the  union  of  but  two  atoms,  the  determination  of  the 
relative  weights  of  the  atoms  would  resolve  itself  into  the  simple 


STANDARD   FOR  ATOMIC   WEIGHTS.  7 

problem  of  determining  the  relative  parts  by  weight  in  which  the 
elements  unite ;  that  this  is  not  the  case,  however,  will  at  once  be 
seen  from  the  fact  that  the  same  elements  can  enter  into  two  or 
more  compounds.  One  thing  must  be  true,  however  —  the  stoichio- 
metric  quantities,  of  necessity,  bear  some  simple  relationship  to  the 
relative  weights  of  the  atoms  themselves ;  and  it  must  be  equally 
true  that,  if  it  is  at  all  possible  actually  to  determine  the  relative 
numbers  which  are  to  be  assigned  to  the  atoms,  such  determinations 
must  be  based  primarily  on  a  discovery  of  the  stoichiometric  quan- 
tities.* The  subsequent  task  is  the  one  of  calling  to  our  aid  all  pos- 
sible guides,  physical  or  chemical  in  nature,  which  will  lead  us  to 
agree  upon  which  ones  of  the  various  multiples,  or  sub-multiples,  of 
the  stoichiometric  quantities  are  to  be  selected  as  expressing  the 
true  relative  weights  of  the  atoms. 

The  methods  by  which  such  an  agreement  has  been  brought 
about,  being  of  a  somewhat  complicated  nature,  are  left  for  discus- 
sion to  a  subsequent  part  of  the  book.t 

The  selection  of  a  standard  by  which  all  other  weights  can  be 
compared  is  as  necessary  in  dealing  with  atoms  as  it  is  in  the  men- 
suration of  distance,  it  being  immaterial  what  standard  is  selected, 
provided  all  of  the  weights  can  easily  and  accurately  be  compared 
with  this.  During  the  first  years  of  our  atomic  hypothesis,  the 
weight  of  the  atom  of  hydrogen,  being  the  smallest  appertaining  to 
any  element,  was  selected  as  unity ;  but  subsequently  this  practice 
was  abandoned  in  favor  of  oxygen,  the  weight  of  the  atom  of  which 
was  placed  at  one  hundred.  Hydrogen  once  more  resumed  its 
original  position  during  the  middle  of  the  century  and,  until 
recently,  all  weights  of  atoms  (atomic  weights),  were  compared  with 
this.  If  we  call  to  our  aid  certain  theories  concerning  the  nature 
of  gases,  a  consideration  which  must  be  deferred  until  the  pupil 
has  become  acquainted  with  a  larger  number  of  chemical  facts,  we 
can  place  the  ratio  between  the  atomic  weights  of  hydrogen  and 

*  The  absolute  weights  of  the  atoms,  being  extremely  small  fractions  of  a 
milligramme,  are  quantities  not  obtainable  with  any  degree  of  accuracy. 

t  As  these  determinations  involved  the  most  painstaking  and  difficult 
manipulations,  as  well  as  very  advanced  views  as  regards  theoretical  deduc- 
tions from  certain  physical  facts,  it  of  necessity  followed  that  wide  differences 
of  opinion,  only  disappearing  within  the  most  recent  times,  were  manifested ; 
indeed,  absolute  certainty  is  not  even  now  attained  or  attainable  as  regards 
the  atomic  weights. 


8 


ATOMIC    WEIGHTS. 


oxygen  at  1  :  15.88,  a  number  which  very  nearly  coincides  with 
1 :  16.  Considerable  uncertainty  exists  as  to  the  accuracy  of  this 
ratio,  for  recent  investigation  has  altered  the  relation  repeatedly. 
If  atomic  weights  are  referred  to  hydrogen  as  unity,  a  recalculation 
of  all  atomic  weights  is  necessary  whenever  investigation  shows 
the  accepted  ratio  between  oxygen  and  hydrogen  to  be  untenable ; 
for  these  constants  have  been  determined  (for  the  greater  number 
of  elements),  directly  or  indirectly,  by  an  investigation  of  com- 
pounds with  oxygen.  It  seems  more  advisable,  therefore,  to  adopt 
\  oxygen  as  a  standard,  and,  so  as  not  to  depart  too  far  from  numbers 
-rendered  familiar  by  accepted  usage,  to  place  the  atomic  weight  of 
this  element  at  16.  By  this  means  the  atomic  weight  of  hydrogen 
becomes  1.008,  a  number  which,  for  all  practical  purposes,  can  be 
placed  at  unity.  If  any  further  correction  in  the  ratio  between 
hydrogen  and  oxygen  becomes  necessary,  such  a  change  will  in- 
volve no  further  calculation.  The  methods  by  which  the  atomic 
weights  have  been  determined  are  not  a  subject  for  discussion  at 
the  present  time ;  indeed,  the  great  majority  of  them  would  be 
entirely  out  of  place  in  an  elementary  treatise ;  suffice  it  to  say, 
that  so  complete  has  been  their  application  that  the  weights  which 
are  placed  in  the  following  table  *  are,  with  unimportant  exceptions, 
accepted  as  correct  by  all  chemists :  — 


NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

ALUMINIUM 

Al 

27 

Erbium  

Er 

166.3 

ANTIMONY    .  . 
ARSENIC.  .  .  . 
BARIUM 

Sb 
As 
Ba 

120. 
75. 
137  43 

Fluorine   .... 
Gadolinium    .  . 
Gallium 

F 

Gd 
Ga 

19. 

156.1 
69. 

Beryllium  .  .  . 
BISMUTH 

Be 
Bi 

9. 

208 

Germanium    .  . 
GOLD  

Ge 

Au 

72.3 
197.3 

BORON 

B 

11. 

HYDROGEN.  .  . 

H 

1.008 

BROMINE 

Br 

79  95 

Indium 

In 

113.7 

Cd 

112 

IODINE           .  . 

I 

126.85 

Caesium. 

Cs 

132  9 

Iridium  

Ir 

193.1 

CAI  CIUM 

Ca 

40 

Fe 

56. 

CARBON  .... 
CERIUM 

C 

Ce 

12. 

140  2 

Lanthanum    .  . 
LEAD 

La 
Pb 

138.2 
206.95 

CHLORINE  .  .  . 
CHROMIUM    .  . 
Cobalt  

Cl 
Cr 
Co 

35.45 
52.1 
59. 

LITHIUM  .... 
MAGNESIUM  .  . 
MANGANESE  .  . 

Li 
Mg 
Mn 

7.02 
24.3 
55. 

Columbium  .  . 
COPPER  .... 

Cb 

Cu 

94. 
63.6 

MERCURY   .  .  . 
Molybdenum  .  . 

Hg 
Mo 

200. 
96. 

*  Compiled  by  F.W.  Clarke;  Jan.  1, 1894.    Jour,  of  the  Am.Chem.Soc.16;  No.  3. 


NUMBER    OF   ELEMENTS. 


NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

Nd 

140  5 

SODIUM 

Na 

23.05 

]N  ickel 

Ni 

58  7 

Strontium    .  .  . 

Sr 

87.6       /// 

NITROGEN.  .  . 
Osmium  .... 
OXYGEN  .... 
Palladium  .  .  . 
PHOSPHORUS  . 
Platinum    .   .  . 
POTASSIUM   .  . 
Praseodymium 

N 
Os 
0 
Pd 
P 
Pt 
K 
Pr 

14.03 
190.8 
16. 
106.6 
31. 
195. 
39.11 
143  5 

SULPHUR.  .  .  . 
Tantalum    .  .  . 
Tellurium    .  .  . 
Terbium    .... 
Thallium  .... 
Thorium   .... 
Thulium   .... 
Tin  

S 
Ta 
Te 
Tb 
Tl 
Th 
Tm 
Sn 

32.06 
182.6 
125.  * 
160. 
204.18 
232.6 
170.7 
119. 

Rhodium    .  .  . 
Rubidium  .  .  . 
Ruthenium    .  . 
Samarium  .  .  . 
Scandium  .  .  . 
Selenium    .  .  . 
SILICON 

Rh 
Rb 
Ru 

Sm 

Sc 
Se 
Si 

103. 
85.5 
101.6 
150. 
44. 
79. 
28  4 

Titanium  .... 
Tungsten  .... 
Uranium  .... 
Vanadium  .  .  . 
Ytterbium   .  .  . 
Yttrium    .... 
ZINC 

Ti 
W 
U 
V 
Yb 
Yt 
Zn 

48. 
184.9 
239.6 
51.4 
173. 
89.1 
65.3 

SILVER    .... 

Ag 

107.92 

Zirconium    .  .  . 

Zr 

90.6 
t 

We  are  acquainted  with  sixty^seven  different  kinds  of  matter, 
no  individual  variety  of  which  has  been  decomposed  into  two 
or  more  simpler  forms ;  but  whether  such  decomposition  will 
ever  occur,  it  is  impossible  to  state.  By  the  union  of  these  elements 
all  substances  known  to  us  are  produced.  By  far  the  greater  pro- 
portion of  matter,  being  composed  of  molecules  containing  two  or 
more  atoms  differing  from  each  other  in  kind,  is  compound  in  its 
nature.  The  individual  atoms  do  not,  except  in  rare  instances, 
exist  as  such;  they  are  united  to  form  molecules,  the  difference 
between  the  molecule  of  the  element  and  that  of  the  compound 
being  that,  while  in  the  former  the  atoms  are  all  of  the  same  kind, 
in  the  latter  they  differ.  Atoms  are  grouped  together  to  form 

*  According  to  Brauner  (Monatshefte  fiirChemie;  10,  445)  tellurium  has 
an  atomic  weight  of  127.6.  Brauner  observes,  however,  that  the  substance 
which  has  heretofore  been  considered  as  an  element  probably  contains  other 
elements  as  well.  The  true  atomic  weight  of  tellurium  he  believes  will  be 
found  to  be  between  125  and  126. 

t  The  ratio  between  the  atomic  weights  of  hydrogen  and  oxygen  is 
1 : 15.88.  The  term  glucinum  is  frequently  used  instead  of  beryllium.  The 
more  important  elements  are  in  capitals.  As  a  matter  of  expediency  the 
pupil  should  memorize  the  atomic  weights  of  a  few  of  the  more  important 
elements.  The  acquirement  of  this  knowledge  is  best  deferred  until  the  indi- 
vidual elements  are  discussed,  when  the  weights  can  be  learned  during  the 
progress  of  the  study. 


10  CHEMICAL   AFFINITY,    CHEMICAL   ENEKGY. 

molecules  which  are  more  or  less  stable,  and  this  stability  of  equi- 
librium must  be  brought  about  by  some  force  acting  between  the 
individual  atoms.  This  force  has  been  compared  to  the  attraction  of 
gravitation,  and  has  by  some  been  considered  to  be  identical  with 
it.  The  attraction  of  gravitation,  however,  is  capable  of  manifesta- 
tion between  bodies  at  a  great  distance  from  each  other,  while  the 
number  of  bodies  acted  on  in  this  manner  by  any  given  body  is 
unlimited.  The  attraction  between  the  atoms  seems  capable  of 
manifestation  only  through  an  extremely  small  interval  of  space, 
and  then  only  between  a  limited  number  of  atoms.  A  new  term  is 
therefore  necessary  to  designate  this  force  which  holds  the  atoms 
in  equilibrium  in  the  molecule,  and,  for  want  of  better  ones,  the 
expressions  "  chemical  affinity,"  or  "  chemism,"  are  used.  Where 
a  very  stable  compound  exists,  the  atoms  composing  it  are  said  to 
have  a  great  affinity  for  each  other.  An  inquiry  into  the  relative 
stability  of  chemical  compounds  is  of  the  greatest  importance  and, 
consequently,  will  frequently  be  made  during  the  progress  of  this 
work ;  but  it  must  not  be  forgotten  that  the  term  "  chemical 
affinity  "  is  used  simply  to  designate  a  force  which  has  never  been 
resolved  into  simpler  factors,  and  of  the  nature  of  which  we  are 
consequently  ignorant. 

When  a  substance,  either  by  reason  of  its  position  or  of  its 
motion,  is  capable  of  performing  work,  it  is  said  to  possess  energy, 
and  work  is  done  when  a  resistance  is  overcome.  The  atoms 
possess  energy  because  they  are,  in  uniting,  capable  of  performing 
work,  by  reason  of  their  chemical  affinity.  Illustrations  of  this 
performance  of  work  by  the  union  of  atoms  are  familiar.  The 
motions  of  machinery  driven  by  steam  can,  with  the  greatest  ease, 
be  traced  to  the  chemical  union  of  the  oxygen  of  the  atmosphere 
with  the  coal  under  the  boilers;  and  the  movements  of  animals 
can  in  the  same  way  be  shown  to  be  derived  from  the  chemical 
energy  of  the  various  substances  which  form  the  nutriment  of  the 
body.  The  energy  possessed  by  the  individual  atoms  can  be 
likened  to  potential  energy  (energy  of  position),  for  by  its  means 
the  atoms  are  capable  of  performing  work,  just  as  is  a  stone  when 
raised  above  the  level  of  the  earth. 

The  measure  of  work  is  the  force  (P),  which  overcomes  resistance,  into 
the  distance  (<S)  through  which  this  force  acts  (L=  P  S).  The  amount  of 
work  which  a  body  is  capable  of  performing,  by  reason  of  its  position,  is  called 


CHEMICAL   ENERGY.  11 

its  potential  energy.  If  the  work  (i)  of  the  force  (P)  is  neutralized  by  a 
negative  force  (P7),  L  is  not  lost,  for  P  can  perform  its  work  (£)  as  soon  as 
P'  ceases  to  act.  A  body  of  the  weight  P,  which  has  been  lifted  through  the 
distance  S,  has  an  amount  of  potential  energy  equal  to  PS,  for  as  soon  as  it 
is  dropped,  it  can  perform  the  work  PS.  The  energy  possessed  by  the  atoms 
is  to  us  different  from  potential  energy,  because  it  has  never  been  resolved 
into  the  factors  PS.  The  amount  of  work  which  a  body  in  motion  is  capable 
of  performing  by  reason  of  that  motion,  is  called  its  kinetic  energy  (energy  of 
motion).  This  is  equal  to  one-half  the  product  of  the  mass  of  the  body  and 

the  square  of  its  velocity  (L  =  — ~ — ).  If  a  part  of  the  potential  energy  pos- 
sessed by  a  body  has  been  used,  and  thereby  a  certain  amount  of  kinetic 
energy  has  been  produced,  then  the  sum  of  the  remaining  capacity  for  work 
and  of  the  produced  kinetic  energy  is  equal  to  the  original  amount  of  potential 
energy;  or,  what  is  the  same  thing,  equal  to  the  kinetic  energy  which  would 
have  been  produced  if  the  entire  store  of  potential  energy  had  been  used  to 
perform  work.  If  the  stone,  which  we  had  supposed  to  be  suspended,  had 
been  allowed  to  fall  to  the  ground,  the  entire  content  of  potential  energy 
would  have  been  changed  to  heat,  which  is  a  form  of  kinetic  energy.  These 
few  remarks  illustrate  the  principle  of  conservation  of  energy,  which  was 
denned  by  Mayer  in  1842. 

What  is  true  of  the  stone  is  also  true  of  the  atoms,  for  their 
potential  energy,  or  better,  chemical  energy  (a  term  which  will  in- 
volve us  in  no  contradictions),  can  be  transformed  into  kinetic 
energy  upon  their  union  to  form  a  compound.  While  we  cannot 
measure  the  amount  of  chemical  energy  possessed  by  the  atoms,  we 
are  able  to  measure  the  kinetic  energy  which  is  produced  when 
they  unite  by  reason  of  their  chemical  affinity,  for  by  far  the 
greatest  amount  of  this  kinetic  energy  is  manifested  as  heat.  If 
the  stone,  which  we  have  used  as  an  illustration,  falls  through 
a  certain  distance,  its  potential  energy  is  converted  into  kinetic 
energy,  and  to  once  more  raise  this  stone  to  its  original  position 
will  require  an  equivalent  expenditure  of  work.  Just  so  with  the 
atoms.  If  two  atoms  have  a  great  chemical  affinity  for  each  other, 
they  possess  a  great  amount  of  chemical  energy,  which  will  be  con- 
verted into  an  equally  great  amount  of  kinetic  energy  when  they 
unite;  and,  in  order  to  separate  the  molecule  so  formed  into  the 
original  atoms,  we  will  have  to  apply  the  same  amount  of  energy 
in  some  form ;  such  a  molecule  is  as  a  consequence  stable,  and  pos- 
sesses an  amount  of  chemical  energy  much  smaller  than  that  of  its 
constituent  atoms.  On  the  other  hand,  two  or  more  atoms  may 
possess  very  little  chemical  affinity  for  each  other,  so  little,  indeed, 


12  METALS    AND   NOT-MET ALS. 

that  an  expenditure  of  energy  is  necessary  to  cause  them  to  unite ; 
therefore,  the  molecule  caused  by  such  a  union  will  have  a  tendency 
readily  to  convert  its  acquired  chemical  energy  into  some  form  of 
kinetic  energy,  and  will  be  unstable.  Molecules  which  have  been 
formed  with  an  evolution  of  heat,  and  which  therefore  possess  less 
energy  than  their  constituent  atoms,  are  said  to  be  the  product  of 
an  exothermic  reaction ;  while  those  which  have  resulted  from  an 
absorption  of  heat  are  the  result  of  an  endothermic  one.  -*The  heat 
given  off  or  absorbed  in  these  changes  has,  in  many  cases,  been 
measured,  and  is  taken  as  an  indication  of  the  chemical  affinity  of 
the  atoms.  The  elements  show  great  variations  in  this  respect; 
some  have  a  great  chemical  affinity  for  each  other,  and  thus  form 
stable  compounds,  others  have  very  little,  while  some  have  none  at 
all,  in  which  case  no  expenditure  of  energy  will  cause  them  to  unite. 
We  must  bear  in  mind  that  any  given  element  varies  much  in  its 
affinity  toward  the  various  other  elements,  and  also  that  what  we 
have  explained  as  regards  compounds  formed  between  individual 
atoms  is  true  as  regards  compounds  formed  from  two  or  more  mole- 
cules. The  various  phases  of  chemical  reaction  which  we  will  sub- 
sequently study,  will, give  abundant  opportunity  for  returning  to 
this  subject. 

The  elements  are  divided  into  two  classes,  the  most  marked 
representatives  of  each  of  which  exhibit  the  sharpest  possible 
chemical  contrast  toward  the  other;  one  of  these  classes  is 
familiar  to  all  of  us :  it  is  that  of  the  metals,  with  the  superficial 
qualities  of  which  substances,  such  as  metallic  lustre,  malleability 
and  ductility,  we  are  tolerably  well  acquainted  ;  in  addition  to  these 
properties,  metals  are  also,  as  a  rule,  good  conductors  of  heat  and 
of  electricity.  At  the  opposite  chemical  extreme  we  find  a  class  of 
elements  which  can  best  be  termed  not-metals.  A  few  of  these  are 
gases,  and  thereby  they  differ  from  the  metals,^  only  one  of  which, 
hydrogen,  exists  in  this  state  at  ordinary  temperatures.  The  not- 
metals,  when  solid,  are  colored  bodies  which  are  brittle,  neither 
malleable  nor  ductile,  and  either  non-conductors  of  electricity  and 
poor  conductors  of  heat,  or  at  least  they  possess  these  two  proper- 
ties in  a  degree  much  inferior  to  that  of  the  metals.  A  number  of 
not-metals,  on  superficial  examination,  will  appear  to  have  metallic 
lustre,  yet  a  closer  inspection  reveals  the  fact  that  this  is  only 
.apparently  the  case,  for,  in  forming  a  thin  section  of  the  element, 


NEGATIVE    AND    POSITIVE   ELEMENTS.  13 

we  find  the  same  to  be  capable  of  transmitting  light.*  Between  the 
two  extremes  we  have  a  considerable  number  of  elements  which 
may  have  the  characteristics  of  both  metals  and  not-metals ; 
examples  of  such  elements  are  arsenic  and  tellurium ;  these  sub- 
stances, while  possessing  a  distinctly  metallic  appearance,  have 
chemical  characteristics,  which  would  cause  them  to  be  more  prop- 
erly classed  as  not-metallic  substances.  The  most  pronounced  metals 
have  a  marked  chemical  affinity  for  the  characteristic  not-metals, 
and  form  stable  compounds  with  them ;  the  less  apparent  the 
contrast  in  the  elements,  the  more  easily  decomposed  will,  in  many 
cases,  the  compound  resulting  from  the  union  of  such  elements  be. 
This  statement  must  not,  by  any  means,  be  taken  as  a  general  rule, 
for  we  find  the  most  stable  substances  resulting  from  the  union  of 
atoms  of  the  same  element,  as  is  the  case  in  the  molecules  of  chlo- 
rine and  of  hydrogen,  and  also  in  molecules  formed  from  elements 
which  closely  resemble  each  other,  as  in  the  compounds  of  sulphur 
and  oxygen.  If  a  compound  formed  of  a  metal  and  a  not-metal  is 
subjected  to  the  action  of  an  electric  current,  the  metal  will  separate 
at  the  negative  pole,  the  not-metal  at  the  positive  one ;  and  for  this 
reason  the  metals  are  called  electro-positive,  while  the  not-metals  are 
electro-negative.  This  rule  was  formerly  supposed  to  be  universal 
in  its  application,  and  a  system  of  chemistry  was  then  established 
which  had  for  its  basis  a  theory  that  all  compounds  are  formed  by 
the  union  of  a  negative  and  a  positive  element  to  form  a  molecule, 
and  that  the  more  complex  substances  are  then  produced  by  the 
union  of  a  negative  and  positive  molecule.  While  the  negative  or 
positive  character  of  the  elements  composing  a  compound  is  undoubt- 
edly of  great  influence  upon  the  nature  of  the  resulting  body,  yet 
a  system  founded  exclusively  on  these  characteristics  has  proved  to 
be  untenable ;  convenience,  however,  causes  us  to  retain  the  expres- 
sions "  negative  "  and  «  positive  "—  the  term  negative  being  used 
to  designate  all  of  the  properties  which  characterize  the  not-metal, 
and  the  term  positive  those  of  the  metal. 

There  are  certain  chemical  characteristics  which  enable  us  to 
draw  the  line  between  metal  and  not-metal  with  tolerable  distinct- 
ness. The  metals  all  form  compounds  with  oxygen,  the  not-metals, 
(with  two  exceptions),  fluorine  and  bromine,  do  the  same ;  these 

*  Some  metals  (gold  foil)  can  imperfectly  transmit  light  when  hammered 
to  the  very  thinnest  possible  leaf. 


14  CHEMICAL   SYMBOLS. 

compounds  are  known  as  oxides.  The  oxides  of  the  metals,  in 
most  cases,  are  classed  as  bases,  a  term  which  is  also  applied  to 
compounds  of  the  metals  with  oxygen  and  hydrogen ;  the  oxides  of 
the  not-metals  are  termed  the  anhydrides  of  acids,  if,  on  addition 
of  water,  they  are  capable  of  forming  acids.  The  contrast  between 
metal  and  not-metal  becomes  apparent  when  these  oxides  interact 
chemically,  for  then  a  new  compound,  a  salt,  is  produced.  Thus 
the  oxide  of  potassium,  a  base,  when  brought  in  contact  with  the 
.oxide  of  sulphur,  an  anhydride  (sulphur  trioxide)  produces  a  salt, 
the  sulphate  of  potassium.  We  shall  more  thoroughly  comprehend 
the  meaning  of  these  terms  as  we  become  better  acquainted  with 
chemical  facts. 

The  writing  of  the  innumerable  chemical  changes  which  take 
place  would  be  complicated  and  a  co-ordination  of  the  phenomena 
rendered  difficult,  were  not  some  system  of  notation  employed 
which  would  express  the  reactions  without  the  trouble  of  writing  in 
full  the  names  and  atomic  weights  of  the  various  elements.  Feeling 
this. need,  some  designation  of  the  elements  by  signs  has  been  em- 
ployed since  the  time  of  the  alchemists.  The  system  introduced 
by  the  Swedish  chemist,  Berzelius,  is  the  one  which  has  been  found 
to  best  answer  all  of  the  requirements,  and  is  in  use  at  the  present 
time.  In  this  the  various  elements  are  designated  by  the  first  letter 
of  their  English  or  Latinized  name  in  capitals,  or,  where  conflict 
would  arise,  by  the  first  letter  of  the  name  in  capitals  followed  by 
some  other  letter,  usually  the  next  following,  in  small  type.  Thus, 
hydrogen  has  the  symbol  H;  oxygen,  0;  sodium,  Na,  from  the 
Latinized  natrium;  mercury,  Hg,  from  the  Latin  hydrargyrum. 
These  symbols  do  not  stand  for  the  visible  elements,  but  represent 
the  atoms,  and  hence  include  the  idea  of  atomic  weights.  Thus  O 
stands  for  an  atom  of  oxygen,  and  means  that,  as  the  atom  of 
oxygen  is  sixteen  times  as  heavy  as  the  atom  of  hydrogen,  we  have 
sixteen  times  as  much  oxygen  by  weight  as  hydrogen,  when  the 
latter  is  expressed  by  the  symbol  H.  By  writing  the  symbols  side 
by  side,  we  express  a  chemical  compound,  as,  for  instance,  H2  0, 
which  stands  for  the  compound  water ;  the  number  2  placed  after 
and  below  the  letter  H  meaning  that  in  one  molecule  of  water 
there  are  two  atoms  of  hydrogen,  the  whole  combination  meaning 
that  in  H2  0  we  have  two  atoms  of  hydrogen  united  to  one  of 
oxygen,  in  the  ratio  of  2  to  16  by  weight ;  H2  0  being  called  the 


CHEMICAL  EQUATIONS.  15 

A 

formula  of  water,  and  the  sum  of  the  atomic  weights  (represented 
by  18),  the  formula  weight.  In  the  case  of  water,  and  in  the  case 
of  many  other  substances  which  can  be  obtained  as  gases,  this 
formula  weight  is  also  the  weight  of  the  molecule,  and  is  therefore 
the  molecular  weight.  The  combination  of  letters,  NaCl,  means 
that  one  atom  of  sodium  is  united  to  one  atom  of  chlorine  to  form 
a  formula  weight  of  the  compound  (sodium  chloride)  ;  at  the  same 
time  it  indicates  that  twenty-three  parts  by  weight  of  sodium  are 
united  to  thirty-five  and  a  half  parts  of  chlorine,  these  numbers 
representing  respectively  the  atomic  weights  of  sodium  and  of 
chlorine.  Whether  NaCl  represents  the  molecular  weight  of 
sodium  'chloride  we  cannot  state,  for  the  substance  may  be  composed 
of  molecules  formed  by  the  union  of  a  number  of  formula  weights, 
Na  01 ;  however,  it  is  extremely  probable  that,  if  such  be  the  case, 
the  molecule  of  sodium  chloride  is  produced  by  the  union  of  a 
number  of  entities  which  all  have  the  composition  N"aCl,  so  that 
there  is  no  great  objection  against  using  the  terms  molecular  weight 
and  formula  weight  interchangeably  in  many  cases.* 

Qur  present  knowledge  seems  to  show  that  the  total  amount  of 
matter  contained  in  the  universe  never  varies.  The  atoms  compos- 
ing chemical  compounds  may  change  their  position  or  manner  of 
grouping,  they  may  be  transferred  from  one  compound  to  another, 
or  the  compound  may  be  decomposed  into  its  elements,  but  the 
atoms  can  neither  be  created  nor  destroyed.  The  entire  science  of 
chemistry  is  based  upon  the  law  of  the  indestructibility  of  matter 
which  was  first  understood  by  Lavoisier,  and  every  chemical  change 
since  his  time  has  but  served  to  prove  its  existence.  As  a  conse- 
quence of  this  law,  the  compounds  produced  by  the  interaction  of 
elements  must  equal  in  weight  the  amounts  taken  before  the 
reaction,  and  the  sum  of  the  weights  of  the  elements  or  compounds 
produced  by  the  decomposition  of  a  compound  substance  must  equal 
the  weight  of  the  substance  originally  employed.  As  our  symbols 
are  used  to  express  atomic  weights  as  well  as  other  characteristics 
of  the  elements,  it  follows  that  we  can  bring  any  chemical  reaction 
into  the  form  of  an  equation ;  thus  H  -J-  Cl  =  H  Cl  means  that 
hydrogen  plus  chlorine  yields  a  compound  of  hydrogen  and  chlorine 
(hydrochloric  acid),  and  it  also  means  that  one  part  by  weight  of 

*  A  number  of  cases  where  the  formula  weight  and  the  molecular  weight 
are  not  identical  will  be  encountered  in  the  subsequent  portions  of  this  work. 


16  CHEMICAL  EQUATIONS. 

hydrogen  plus  35.5  parts  by  weight  of  chlorine  (the  latter  number 
being  the  atomic  weight  of  that  element),  equal  36.5  parts  by 
weight  of  hydrochloric  acid. 

The  formula  HgO  =  Hg  -j-  0  indicates  that  the  oxide  of  mer- 
cury is  decomposed  into  mercury  and  oxygen,  and  that  the  weight 
of  the  oxide  of  mercury  used  is  equal  to  the  sum  of  the  weights  of 
oxygen  and  mercury  produced. 

C204H2  =  C02  +  CO+H20 

means  that  a  compound  of  carbon,  hydrogen,  and  oxygen  (oxalic 
acid)  can  be  decomposed  into  C02  (carbon  dioxide),  CO  (carbon 
monoxide),  and  H2  0  (water),  and  that  the  weight  of  the  oxalic  acid 
taken  is  equal  to  the  sum  of  the  weights  of  carbon  dioxide,  carbon 
monoxide,  and  water  produced.  The  atomic  weights  of  carbon, 
oxygen,  and  hydrogen  are  12,  16  and  1  respectively ;  hence 
(2  X  12)  -f  (4  X  16)  -f-  2  =  90  parts  by  weight  (of  oxalic  acid) 
yield,  44  -f-  28  +  18  =  90  parts  by  weight  of  carbon  dioxide,  carbon 
monoxide,  and  water. 

Some  doubts  will  arise  in  the  beginner's  mind  as  to  the  legiti- 
macy of  the  numbers  used  as  atomic  weights  ;  he  will  naturally  ask, 
why  is  the  atomic  weight  of  carbon  12  and  not  6  ?  Why,  for 
instance,  should  we  suppose  the  formula  of  carbon  dioxide  to  be 
C02  and  not  C2  02  ?  These  very  doubts  were  entertained  by  chem- 
ists during  the  first  six  decades  of  the  century,  and  only  by  pains- 
taking investigation,  with  the  aid  of  a  number  of  physical  laws, 
have  they  been  removed.  At  the  present  period  of  our  study  we 
must  accept  the  atomic  weights  as  they  are  given ;  a  discussion  of 
the  various  reasons  for  this  acceptance  would  be  far  beyond  the 
scope  of  an  elementary  treatise ;  but  such  as  can  be  understood  will 
be  introduced  at  a  subsequent  time.  Chemists,  as  a  rule,  are  scep- 
tical to  a  degree,  and  the  fact  that  our  present  atomic  weights  are 
believed  in  by  the  great  majority  must  be  a  sufficient  guarantee  for 
their  accuracy.  The  pupil  must  remember  that  chemical  symbols 
stand  for  material  substances ;  he  must  not  be  led  to  consider  the 
element  nitrogen,  for  instance,  as  being  simply  the  symbol  N,  nor 
potassium,  the  symbol  K,  but  should  always  bear  in  mind  that  these 
letters  stand  for  the  atoms  of  existing  elements,  endowed  with  various 
properties,  which  elements  and  properties  would  be  in  existence 
even  if  no  system  of  chemical  notation  had  ever  been  established. 


ARRANGEMENT  OF  THE  ELEMENTS.  17 

The  elements  are  not  substances  any  one  of  which  differs  in 
properties  from  all  of  the  others;  indeed,  no  one  element  exists 
which  does  not  exhibit  characteristics  which  mark  its  resemblance 
to  a  number  of  its  fellows.  Roughly  speaking,  the  elements  are 
divided  into  metals  and  not-metals,  the  metals  all  bearing  a  certain 
resemblance  to  each  other,  and  the  not-metals  all  being  more  or  less 
alike ;  yet  both  metals  and  not-metals  are  naturally  divided  into 
groups,  of  which  the  individual  members  have  a  strong  family  like- 
ness, while  these  various  groups  also  have  a  number  of  points  in 
common.  We  shall  see,  at  a  subsequent  time,  that,  by  arranging 
the  elements  in  the  order  of  their  atomic  weights,  beginning  with 
the  one  which  has  the  smallest  number,  this  resemblance  will 
become  most  apparent.  By  constructing  a  table  of  the  elements 
arranged  in  this  manner,  we  have  produced  a  system  of  classifica- 
tion, a  careful  study  of  which  has  shown  us  that  the  characteristics 
of  any  element  are  determined  by  the  atomic  weight  of  that 
element ;  so  that  if  we  are  thoroughly  familiar  with  this  arrange- 
ment, we  can  describe  the  properties  of  an  element  with  tolerable 
accuracy,  without  any  previous  knowledge  of  its  chemical  deport- 
ment, provided  that  we  are  acquainted  with  its  atomic  weight. 
This  system  will  provide  the  order  in  which  the  various  elements 
will  be  discussed. 

Having  by  this  brief  introduction  learned  some  of  the  funda- 
mental theories  upon  which  our  science  is  based,  we  will  now  go  on 
to  the  special  descriptive  portion  of  chemistry.  In  so  doing  we 
shall  find  it  expedient  to  become  acquainted  first  with  a  typical  note 
metal,  then  with  a  metal,  and  immediately  following  this  with  the 
compound  produced  by  the  chemical  union  of  these  two  contrasting 
elements.  The  not-metal  will  be  oxygen,  the  metal,  hydrogen,  and 
the  compound,  water. 


18  OXYGEN;    OCCURRENCE,   HISTORY. 


CHAPTER  II. 

OXYGEN. 

Symbol,  0 ;  atomic  weight,  16 ;  specific  gravity,  air  =  1,  is  1.10535, 
H2  =  2,  is  31.76  ;  1  c.c.  0  at  0°  and  .76m.  pressure  weighs 
.00142952  gram. 

OXYGEN  is  the  element  which  occurs  in  greatest  quantity  upon 
our  planet ;  it  forms  about  47.3  per  cent  of  the  solid  portion  of  the 
earth,  85.8  per  cent  of  the  ocean,  and,  including  its  occurrence  in 
the  atmosphere,  about  50  per  cent  of  the  total  substance  of  the 
globe.  It  is  present  in  the  atmosphere  as  uncombined  oxygen, 
mixed  with  nitrogen  and  a  few  other  gases,  and  constitutes  one- 
fifth  of  the  entire  gaseous  envelope  of  the  earth.  United  with 
hydrogen  it  produces  water,  forming  eight-ninths  of  that  substance  ; 
by  far  the  greater  portion  of  the  crystalline  rocks  contain  oxygen 
combined  with  other  elements ;  the  soil  and  the  various  forms  of 
vegetable  and  animal  life  contain  this  element  chemically  united, 
and  also  to  some  extent  as  free  oxygen. 

Oxygen  was  isolated  on  August  1st,  1774,  by  an  English  chem- 
ist, Joseph  Priestley,  who  prepared  it  by  heating  "  red  precipitate  " 
(red  oxide  of  mercury)  in  the  focus  of  a  burning  glass  exposed  to 
the  sun's  rays.  He  termed  the  gas  dephlogisticated  air  because, 
according  to  the  theory  then  held,  a  substance,  in  burning,  gave  off 
an  hypothetical  principle  called  phlogiston,  which  was  subsequently 
thought  to  be  identical  with  hydrogen.  As  oxygen  is  capable  of 
supporting  combustion  in  a  very  energetic  manner,  it  was  supposed 
that  the  element  must  necessarily  be  able  to  take  the  phlogiston 
from  the  burning  substance;  and  from  this  property  the  gas 
received  its  first  name. 

In  order  to  prepare  oxygen,  the  following  methods  can  be 
resorted  to :  — 

Red  oxide  of  mercury  is  heated  in  a  hard  glass  tube ; !  *    it  then 

.*  The  superior  numbers  in  the  text  refer  to  the  numbers  of  the  notes  in 
the  laboratory  appendix. 


OXYGEN;  PREPARATION. 


19 


breaks  down  into  mercury  and  oxygen.     The  formula  of  this  oxide 
is  Hg  0,  which  means  that  in  one  formula  weight  of  mercuric  oxide 


]D p  ID ]p  ID 


Fig.   1. 

there  are  united  one  atom  of  mercury  and  one  of  oxygen.  The  re- 
action is  represented  as  follows  :  —  Hg  0  =  Hg  -|-  0.  The  atomic 
weight  of  mercury  is  200,  that  of  oxygen  16 ;  hence  the  equation 
indicates  that  216  parts  by  weight  of  the  oxide  of  mercury  yield 
200  parts  of  mercury  and  16  parts  of  oxygen  upon  heating. 

Black  oxide  of  manganese  (manganese  dioxide)2  is  heated  in  an 
iron  tube,  the  oxygen  so  prepared  being  collected  over  water.3 

The  reaction  is  as  follows: — 

3  Mn  02  =  Mn3  04  -f-  2  0  ;  the  substance  formed  is  an  oxide  of 
manganese  containing  less  oxygen,  in  proportion  to  the  quantity 
of  manganese,  than  does  Mn  02 .  As  one  formula  weight  of  Mn3  04 
contains  three  atoms  of  manganese,  it  follows  that,  in  writing  the 
equation,  three  times  one  formula  weight  of  Mn  02  must  be  used, 
and  this  is  indicated  by  placing  the  coefficient  3  before  the  formula 
Mn  02 .  As  the  atomic  weight  of  manganese  is  55,  it  follows  that 
261  parts  by  weight  of  manganese  dioxide  yield  229  parts  of  the 
oxide  containing  less  oxygen  and  32  parts  of  oxygen.  From  the 
examples  cited,  the  meaning  of  chemical  equations  will  be  suf- 
ficiently understood,  so  that  the  student,  by  consulting  the  table  of 
atomic  weights,  will  be  able  to  determine  the  relations  by  weight 
in  all  others.  We  shall  therefore,  in  the  future,  be  contented  with 


20 


OXYGEN;  .PREPARATION. 


writing  equations  without  indicating  the   quantities  of   the  sub- 
stances reacting.* 

Xeither  of  the  methods  which  have  been  given  is  of 
great  practical  use  in  the  preparation  of  oxygen.     Better 
results  are  obtained  by  heating  chlorate  of  potassium  (a 
white,  crystalline  substance,  somewhat  resembling  com- 
mon salt4)  in  a  flask  connected  with  a 
delivery  tube,  as  in  Fig.  2.     Chlorate 
of  potassium  breaks  down  as  follows  : 


0 


KC103  =  KCl-f  30.f 


(Chlorate  of  potassium  yields  potas- 
sium chloride  and  oxy- 
gen.) The  most  approved 
method  of  preparing  oxy- 
gen for  laboratory  use, 
is  by  heating  a  mixture 
of  chlorate  of  potassium 
and  manganese  dioxide, 
the  apparatus  used  being  F<9-  2- 

the  same  as  for  that  of  the  previous  experiment.5  By  this 
means  the  oxygen  passes  off  at  a  much  lower  temperature,  and  quite 
regularly.  The  part  which  the  manganese  dioxide  plays  in  the  re- 
action is  not  as  yet  definitely  understood,  for  it  remains  as  manga- 
nese dioxide  in  the  final  result.  It  is  not  improbable  that  the 
substance  takes  up  oxygen  from  the  chlorate  of  potassium,  thus 
forming  an  oxide  of  manganese  containing  relatively  more  oxygen 
than  does  Mn02  (a  so-called  higher  oxide),  which  then  readily 
breaks  down  into  manganese  dioxide  and  oxygen,  the  black  oxide 
of  manganese  thus  acting  as  a  conveyer  of  oxygen.  Substances 

*  It  is  just  as  well  for  beginning  students  to  neglect  equations,  and  to  fix 
their  attention  more  especially  on  learning  the  facts,  and  also  upon  familiariz- 
ing themselves  with  the  appearance  and  character  of  the  substances  used  in 
the  reactions.  The  fact  that  these  reactions  take  place  only  between  definite 
masses  of  matter  must,  however,  never  be  lost  sight  of. 

t  The  reaction,  as  represented,  gives  the  final  result  obtained  by  heating 
chlorate  -of  potassium.  There  is  in  reality  an  intermediary  product  formed. 
For  a  full  description  of  the  change  see  Chapter  xviii.  The  pupil  should,  by 
means  of  the  table  of  atomic  weights,  determine  the  relationships  by  weight 
in  this  and  in  a  large  number  of  subsequent  equations. 


OXYGEN;  PROPERTIES,  OXIDES.  21 

which  act  in  this  manner,  their  influence  not  being  clearly  under- 
stood, are  said  to  act  by  catalysis.  All  of  the  methods  for  prepar- 
ing oxygen  which  have  been  cited  have  one  thing  in  common ;  the 
substances  decomposed  contain  oxygen,  which  is  driven  off  by 
heat,  there  remaining  either  elements  or  compounds  containing  pro- 
portionally less  oxygen  than  the  original  material.  As  energy  is 
added  to  effect  the  separation,  the  substances  produced  contain  a 
greater  amount  of  chemical  energy  than  those  employed;  while  all 
of  these  reactions  belong  to  the  class  termed  chemical  analyses.* 

Oxygen  is  colorless,  odorless,  and  tasteless ;  it  is  converted  to  a 
liquid  by  a  pressure  of  50  atmospheres  at  a  temperature  of  —  118°. 
A  certain  temperature  exists,  constant  for  any  given  gas,  but  vary- 
ing with  different  gases,  above  which  they  cannot  be  liquefied  by 
any  pressure.  This  temperature  is  called  the  critical  temperature ; 
for  oxygen  it  is  at  —  118°.  Liquid  oxygen  boils  at  —  182°  at  740 
mm.  pressure,  and  at  —  197°  at  25  mm.f  Oxygen  is  but  slightly 
soluble  in  water,  100  volumes  of  that  liquid  dissolving  4.1  volumes 
of  oxygen  at  5°.  $  Small  as  this  amount  is,  it  is  sufficient  to  fur- 
nish the  material  required  by  fishes  for  their  physiological  functions, 
and  this  solubility  is  therefore  of  the  highest  importance;  water 
containing  no  free  oxygen  is  incapable  of  supporting  marine  life. 

Oxygen  forms  chemical  compounds  (called  oxides)  with  all  other 
elements  excepting  fluorine  and  bromine,  and  it  will  combine  with 
the  latter  element  provided  some  metal  is  also  a  constituent  of  the 
compound.  In  the  formation  of  many  of  the  oxides,  a  large  amount 
of  heat  is  produced ;  this  may  be  so  great  that  the  substance, 
whether  a  burning  solid  or  a  vapor,  will  become  incandescent ;  this 
phenomenon  takes  place  with  greater  intensity  in  oxygen  than  in 
the  air.  That  the  latter  statement  is  true  may  be  proved  by  placing 
a  glowing  pine  chip  in  a  jar  of  oxygen,  for  then  the  wood  will 
at  once  burst  into  flame. 

The  phenomena  attendant  upon  the  burning  of  bodies  in  oxygen 

*  Reactions  in  which  a  compound  body  is  decomposed  into  two  or  more 
simpler  ones  are  termed  chemical  analyses;  those  in  which  a  more  complex 
body  is  formed  from  two  or  more  simpler  ones  are  termed  chemical 
syntheses. 

t  Dewar  and  Fleming,  Philosoph.  Mag. ;  34,  326. 

t  By  this  expression  is  meant  that  100  liters  of  water  would  dissolve  4.1 
liters  of  the  gas. 


22  OXYGEN;   COMBUSTION. 

are  all  essentially  alike,  so  that  a  very  few  examples  will  serve  to 
illustrate  their  character :  — 

If  a  piece  of  phosphorus  is  kindled  and  then  placed  in  a  jar  of  the  gas,  it 
will  continue  to  burn  with  a  dazzling  white  light,6  forming  a  dense  smoke.  This 
smoke  consists  of  fine  particles  of  the  solid  oxide  of  phosphorus,  produced  by 
the  combustion;") this  oxide  contains  eighty  parts  by  weight  of  oxygen  for  every 
sixty-two  parts  by  weight  of  phosphorus ;  its  formula  is  therefore  P2O5  (see 
table  of  atomic  weights).  The  action  can  be  represented  as  follows:  — 

2  P+5  O=P2O5. 

Sulphur,  which  has  been  ignited  in  the  air,  will  continue  to  burn  in  oxygen 
with  a  brilliant  blue  flame ; 6  in  this  case  the  product  of  the  combination  is  a 
gas,  composed  of  equal  parts  by  weight  of  sulphur  and  of  oxygen ;  in  order  to 
have  the  foregoing  composition,  this  oxide  of  sulphur  must  contain  two  atoms 
of  oxygen  to  one  of  sulphur;  it  is  termed  sulphur  dioxide. 

S+20=SO,. 

A  piece  of  glowing  carbon  becomes  brightly  incandescent  when  placed  in 
oxygen,  the  substance  burning  to  form  a  gaseous  compound,  carbon  dioxide, 
(C02):-  ^,0^ 

C  +  20  =  C02. 

The  three  oxides  produced  by  the  above  reactions  are  oxides  of 
not-metals,  and  hence  bear  the  characteristics  of  the  anhydrides 
of  acids.  (See  page  14.) 

Even  substances  such  as  iron,  which  we  do  not  ordinarily  consider  as 
being  combustible,  burn  readily  in  oxygen.  If  a  steel  watch  spring  is  heated 7 
and  placed  in  a  jar  of  the  gas,  the  substance  will  continue  to  burn.  The  heat 
produced  is  sufficient  to  cause  the  iron  to  burn,  while  sparks  of  white  hot 
iron  oxide  spatter  around  the  jar  in  which  combustion  is  taking  place; 
these  particles  may  even  be  hot  enough  to  become  fused  into  the  sides  of 
the  jar. 

3  Fe+4  O  =  Fe3O4* 

From  these  few  examples  we  can  form  some  idea  of  the  readi- 
ness with  which  oxides  are  formed ;  and  what  is  true  of  the  elements 
referred  to  also  holds  good  with  a  large  number  of  other  ones. 
When  a  burning  substance  is  a  gas,  or  when,  in  burning,  it  either 
gives  off  a  gas  or  is  converted  into  a  gas,  then  a  flame  is  produced 
(such  is  the  case  in  the  burning  of  phosphorus  and  of  sulphur) ;  on 
the  other  hand,  when  no  vapor  is  present  to  be  heated  to  incandes- 

*  The  oxide  Fe3O4 ,  occurring  as  a  mineral  called  magnetite,  is  the  oxide 
of  iron  always  formed  at  high  temperatures  when  the  metal  is  in  the  presence 
of  oxygen. 


OXYGEN;    KINDLING   TEMPERATURE.  23 

cence,  then  the  burning  substance  simply  glows  (this  is  seen  in  the 
combustion  of  carbon).^  In  a  restricted  sense,  the  term  combustion 
refers  only  to  the  unidh  of  various  substances  with  oxygen,  with 
the  evolution  of  light  and  heat.  During  any  chemical  process  in 
which  two  elements  directly  unite,  heat  is  liberated;  but,  if  the 
union  of  a  body  with  oxygen  takes  place  slowly,  so  that  in  any 
given  interval  of  time  only  very  small  amounts  of  the  substances 
under  consideration  were  to  combine,  we  might  not  be  able  to  note 
any  increase  of  temperature,  because  the  inconsiderable  quantity  of 
heat  given  off  during  this  period  would  be  conducted  away  by  the 
surroundings.  The  rusting  of  iron  consists  of  such  an  oxidation, 
without  any  perceptible  rise  in  temperature ;  it  is,  nevertheless,  just 
as  much  a  form  of  combustion  as  the  more  rapid  and  brilliant  one 
referred  to  above.  Combustion  may,  therefore,  take  place  either 
slowly  or  rapidly.  A  substance  must  be  heated  to  a  certain  point 
before  slow  can  be  converted  into  rapid  combustion,  and  this  tem- 
perature is  called  the  kindling  temperature.  The  kindling  temper- 
ature varies  for  different  substances,  but  is  constant  for  any  given 
substance,  and  if  the  burning  body  is  cooled  below  that  point,  an 
extinction  of  the  fire  results.  AVe  have  daily  evidence  of  this 
when  lamp  flames  are  blown  out,  for  this  process  consists  simply  in 
cooling,  below  their  kindling; temperature,  the  gases  which  form  the 
flame ;  while,  on  the  other  hand,  the  application  of  a  lighted  match 
to  a  gas  burner  serves  only  to  heat  the  escaping  gases  to  their  kind- 
ling temperature.  If  a  substance  is  undergoing  slow  combustion,  in 
such  a  situation  that  the  heat  given  off  cannot  be  readily  conducted 
away,  the  temperature  of  the  whole  will  gradually  rise  until  the 
kindling  point  is  reached ;  and  in  this  manner  spontaneous  combus- 
tion takes  place  in  heaps  of  oiled  rags,  where  the  oil  is  being  slowly 
oxidized  by  the  oxygen  in  the  atmosphere.  Substances  which  are 
capable  of  oxidation  possess  a  certain  amount  of  chemical  energy 
in  the  presence  of  oxygen,  and,  when  they  unite  with  that  element, 
this  energy  is  converted  into  heat ;  but,  as  the  amount  of  energy  in 
any  given  body  must  be  constant  before  combustion,  no  matter 
whether  the  process  is  to  be  slow  or  rapid,  therefore,  provided  the 
products  of  combustion  do  not  vary,  it  follows  that  the  amount  of 
heat  liberated  must  be  the  same  in  either  case.  It  sometimes 
occurs,  however,  that  the  substances  formed  by  slow  combustion 
may  be  chemically  different  from  those  by  rapid,  when,  of  course, 


24  OXYGEN;    OXIDES. 

the  amount  of  heat  liberated  woiild  depend  upon  the  product 
formed.* 

It  would  be  erroneous  to  suppose  that  the  phenomena  of  combus- 
tion are  apparent  only  when  oxygen  is  taking  part  in  the  perform- 
ance; for,  as  the  process  of  direct  chemical  union  between  two  bodies 
is  always  accompanied  by  heat,  it  follows  that  in  the  union  of  other 
elements  or  compounds  this  may  be  so  great  that  the  substances  re- 
acting begin  to  glow  or  burst  into  flame.  Thus,  for  instance,  hydro- 
gen can  burn  in  chlorine  just  as  readily  as  it  can  in  oxygen;  so 
that,  in  the  very  broadest  sense,  combustion  would  refer  to  the 
union  of  any  substances  with  the  evolution  of  heat.  This  very 
broadening  of  the  sense,  however,  would  rob  it  of  its  significance,  so 
that  the  term  combustion  might,  perhaps,  more  properly  refer  only 
to  the  union  of  any  substance  with  oxygen,  and  with  the  evolution 
of  light  and  heat;  while,  instead  of  slow  combustion,  the  term  slow 
oxidation,  or,  in  the  case  of  other  elements,  some  more  particular 
term,  would  be  preferable. 

Two  elements  which  are  capable  of  directly  uniting  must  possess 
a  certain  amount  of  chemical  energy,  just  as  a  stone  raised  above 
the  ground  possesses  potential  energy ;  and  the  chemical  energy  is 
converted  into  heat  when  they  combine,  just  as  the  potential  energy 
of  the  stone  is  changed  to  kinetic  on  falling.  Now,  after  the  stone 
has  fallen,  it  requires  just  as  much  energy  to  bring  it  once  more 
back  to  its  original  position  as  was  given  off  by  it  in  its  descent. 
Just  so  with  chemical  compounds.  In  order  to  decompose  them,  as 
much  energy  must  be  added  as  was  given  off  in  their  formation. 
Therefore,  in  decomposing  a  body  formed  by  combustion  or  slow 
oxidation,  just  as  much  heat,  or  other  form  of  energy,  must  be 
applied  as  was  given  off  in  the  production  of  the  compound.  The 
heat  (or  the  quantity  of  any  form  of  energy)  necessary  for  the 
decomposition  of  a  substance,  is  therefore  equal  to  the  heat  (or 
energy)  given  off  in  its  formation ;  it  follows  that  those  bodies  which 
burn  energetically  in  oxygen,  with  the  evolution  of  much  heat  and 
light,  will  necessarily  form  stable  oxides. 

The  bodies  produced  by  the  union  of  the  various  elements  with 
oxygen  are  termed  oxides,  and  we  are  acquainted  with  the  oxides 
of  all  elements  excepting  those  of  fluorine  and  bromine.  The  great- 

*  For  example,  phosphorus,  when  slowly  oxidized,  forms  phosphorus 
trioxide;  when  hurned  it  forms  phosphorus  pentoxide. 


OXYGEN;   FORMATION   OF   OXY-SALTS.  25 

est  diversity  of  characteristics  exists  among  these  compounds,  but, 
as  we  have  seen  in  the  introduction,  they  can  in  most  instances 
be  classed  either  as  bases  or  as  anhydrides  of  acids;  the  bases 
are  oxides  of  metals  which  exhibit  the  same  contrast  toward  the 
oxides  of  not-metals  (called  anhydrides)  as  the  metal  itself  does 
toward  the  not-metal,  with  this  difference,  however,  that  the  con- 
trast is  intensified  by  the  addition  of  oxygen  to  the  not-metal  in 
the  formation  of  the  anhydride.  The  result  of  the  chemical  union 
of  base  and  anhydride  is  a  salt  (usually  possessing  somewhat  indif- 
ferent chemical  properties),  and  the  product  of  the  union  of  metal 
with  not-metal  not  infrequently  has  the  characteristics  of  a  salt', 
such  an  instance  is  found  in  common  salt  (sodium  chloride,  Na  Cl). 
Examples  of  the  oxides  of  metals  which  are  bases  are :  — 

Zinc  oxide,  Zn  0 ;  sodium  oxide,  Na2  0 ;  magnesium  oxide, 
MgO. 

Examples  of  the  anhydrides  of  acids  are :  — 

Sulphur  trioxide,  S03 ;  phosphorus  pentoxide,  P2  05 ;  carbon 
dioxide,  C02. 

The  production  of  salts  from  the  union  of  these  two  contrasting 
substances  can  be  represented  as  follows :  — 

ZnO  +  S03  =ZnS04, 

Zinc  oxide  +  Sulphuric  anhydride    =  Zinc  sulphate. 

3  Na2  0  +  Pr05  =  2  Na3  P04 , 

Sodium  oxide         +  Phosphoric  anhydride  =  Sodium  phosphate. 

MgO  +C02  =MgC03, 

Magnesium  oxide  +  Carbonic  anhydride     =  Magnesium  carbonate ; 

and  if  we  compare  these  reactions  with  the  formation  of  common 
salt  from  sodium  and  chlorine :  — 

Na       +  Cl         =  Na  Cl, 

Sodium  +  Chlorine  =  Sodium  Chloride, 

the  similarity  in  the  formation  of  salts  from  the  two  contrasting 
compounds  and  from  two  contrasting  elements  is  quite  evident. 

There  are,  in  addition  to  the  oxides  referred  to  above,  a  large 
number  which  are  neither  basic  nor  acidic.  To  this  class  belong 
the  oxides  known  as  hyperoxides,  of  which  manganese  dioxide  is  an 
example.  All  of  these  have  the  metallic  character  of  the  metal 
about  neutralized  by  the  not-metallic  oxygen,  so  that  where  they 
are  capable  of  taking  up  more  oxygen,  upon  such  addition  the 


26  OXIDES  ;   NOMENCLATIVE. 

resulting  compound  bears  the  characteristics  of  the  negative  an- 
hydride of  an  acid;  and  when  they  lose  oxygen  to  form  a  lower 
oxide,  the  resulting  substance  has  more  or  less  the  characteristics  of 
a  base.  They  are  on  the  turning  point  from  base  to  anhydride, 
and  all  of  them  lose  a  part  of  their  oxygen  readily,  as  does  manga- 
nese dioxide.  A  number  of  the  oxides  of  the  not-metals,  also,  are 
not  anhydrides  of  acids,  and  hence  do  not  directly  unite  with  the 
bases  to  form  salts.  All  of  these  oxides  contain  but  little  oxygen, 
as,  for  example,  carbon  monoxide,  CO ;  nitrous  oxide,  N2  0. 

Many  of  the  elements  are  capable  of  forming  a  number  of 
oxides  (nitrogen  can  even  form  five),  and  these  vary  in  character 
with  the  increase  in  the  relative  quantity  of  oxygen.  A  number  of 
metals  can  form  two  oxides,  both  of  which  are  bases,  and  in  such  a 
case  the  oxide  containing  the  lesser  amount  of  oxygen  relatively  to 
the  quantity  of  metal  is  designated  by  the  suffix  ous  attached  to  the 
name  of  the  metal,  the  one  containing  the  greater  by  ic ;  thus, 
copper  forms  two  oxides,  cuprous  oxide,  Cu2  0  ;  cupric  oxide,  Cu  0  ; 
iron  has  the  same  property,  and  the  two  oxides,  Fe  0  and  Fe2  03 
are  termed  ferous  and  ferric  oxides  respectively ;  mercury  can 
form  mercurous  and  mercuric  oxides  (Hg2  0  and  Hg  0).  The  salts 
derived  by  the  union  of  these  oxides  with  the  anhydrides  of  acids 
are  designated  in  a  similar  manner ;  for  example  :  — 

Ferrous  sulphate,  Fe  S  04 ,  is  produced  by  the  union  of  ferrous 
oxide  with  the  anhydride  of  sulphuric  acid :  — 

Fe  0  +  S03  =  Fe  S04 

and  ferric  sulphate,  Fe2  (S  04)3 ,  is  formed  by  uniting  ferric  oxide 
with  sulphur  trioxide :  — 

Fe203+3S03=Fe2  (S04)3. 

The  suffix  ous  indicates,  when  attached  to  the  name  of  the  metallic 
element  in  a  salt,  that  this  salt  is  derived  from  an  oxide  containing 
less  oxygen,  in  proportion  to  a  fixed  amount  of  the  metal,  than  does 
some  other  basic  oxide  of  the  same  element. 


HYDROGEN;   OCCURRENCE,    HISTORY.  27 


CHAPTER   III. 

HYDROGEN. 

Symbol,  H  ;  atomic  weight,  1.008  ;  specific  gravity,  air  =  1,  is  .06960  ; 
1  c.c.  H  weighs  .00009001  gram  at  0°  and  .76  meters  pressure. 

FREE  hydrogen  is  found  in  nature  only  in  small  quantities ;  it 
occurs  in  the  gases  which  escape  from  petroleum  wells,  in  natural 
gas,  and  in  the  exhalations  of  some  volcanoes.  It  is  said  to  have 
been  found  in  a  condensed  state  in  some  meteorites,  and  it  is  fre- 
quently given  off  in  processes  of  fermentation  and  decay.  The 
element  is,  however,  contained  in  enormous  quantities  in  the  chro- 
mosphere of  the  sun,  the  protuberances  observed  during  eclipses 
consisting,  for  the  greater  part,  of  hydrogen  which  is  subjected  to 
intense  local  disturbances.  The  presence  of  so  much  free  hydrogen 
in  the  atmosphere  of  the  sun  is  accounted  for  by  the  fact  that  the 
temperature  of  that  body  is  so  high  that  it  renders  all  chemical 
union  impossible.  The  fixed  stars,  Sirius  for  example,  also  contain 
large  quantities  of  uncombined  hydrogen ;  on  the  earth  the  total 
amount  of  free  and  combined  hydrogen  is  only  about  1  per  cent  of 
the  entire  mass. 

Hydrogen  was  first  described  as  a  peculiar  form  of  air  by 
Cavendish  in  1766,  although  it  had  previously  been  observed  by 
Paracelsus,  Boyle,  and  others,  who  obtained  it  by  the  action  of 
dilute  acids  on  certain  metals,  but  confused  it  with  other  combus- 
tible gases.*  By  Cavendish,  Priestley,  and  contemporaneous  chem- 
ists, it  was  at  one  time  considered  to  be  pure  phlogiston,  the  ele- 
ment which  was  supposed  to  be  given  off  by  substances  in  burning. 
Lavoisier  explained  the  composition  of  water  at'  a  later  date,  and 
proved  that  substance  to  be  an  oxide  of  hydrogen ;  by  this  means  he 
demonstrated  that  hydrogen  is  a  substance  which  can  act  like  a 
metal  when  it  forms  an  oxide,  and  that  therefore  it  could  not  be 
identical  with  phlogiston.  Lavoisier  gave  the  element  the  name  of 

*  See  Kopp,  Geschiclite  der  Chemie,  3,  260. 


28 


HYDROGEN ;    PREPARATION. 


hydrogene  (from  i>8<ap  water  and  the  root  yei/  to  produce),  as  water 
is  produced  by  burning  the  gas  in  air  or  oxygen. 

The  methods  for  preparing  hydrogen  are  as  follows :  — 

1.  By   the   decomposition    of  water   by   means    of    the    electric 
current,  when  hydrogen  separates  at  the  negative  pole  and  oxygen 
at  the  positive  one.     The  apparatus  used  in  performing  this  experi- 
ment will  be  seen  in  Fig.  3.*  f 

2.  By  decomposition    of  water  by  means  of  metals. 
The  most  pronouncedly  metallic  elements,  such  as  so- 
dium   and    potassium,   have    such   a   great    affinity  for 
oxygen  that  they  will  unite  with  the 

oxygen  combined  in  water,  and  expel 
the  hydrogen  even  at  ordinary  temper- 
atures.    If  a  piece  of  sodium,  the  size 
of  a  large  pea,  is  placed   in  water,   a 
violent  reaction  8  takes  place,  the  heat 
evolved  melts  the  sodium,  which,  being 
specifically  lighter  than  water,  floats 
on  the  surface ;  hydrogen  is  at  the 
same  time    set   at   liberty.     If   the 
water  is  thickened  by  starch  paste, 
so  that  the  molten  metal  cannot  move 
freely  on  its  surface,  the  hydrogen 
will  be  heated  to  its  kindling  tem- 
perature and  take  fire.    The  gas  can 
be  collected  by  placing  the 
sodium  in  a  small  wire  cage 
which  is  placed   under  the 
surface   of   the  water   in   a 
pneumatic    trough,    and   by 
then  inverting  a  test-tube  filled  with  water  over  the  metal,  bubbles 
of  hydrogen  arise  and  soon  fill  the  tube.9     (Fig.  4.)     The  water 

*  Pure  water  is  a  not-conductor  of  electricity,  or  at  least  very  nearly  a 
not-conductor,  and  hence  a  little  sulphuric  acid  must  be  added  to  the  water 
before  performing  the  experiment. 

t  It  will  be  noticed  during  the  experiment  that  for  each  cubic  centimeter 
of  oxygen  liberated  at  the  positive  pole,  we  have  2  c.c.  of  hydrogen  formed  at 
the  negative. 


Fig.  3. 


HYDHOGEN  :    PREPARATION. 


29 


Fig.  4. 


in  which  the  sodium  has  been  dissolved  has  acquired  a  soapy  feel 
and  an  alkaline  taste ;  it  contains  sodium  hydroxide. 

Na        +  HOH  =  Na  OH  -f  H. 

Sodium  +  water    =  Sodium  hydroxide  +  hydrogen. 

A  reaction  of  this  kind  is 
called  one  of  substitution, 
for  the  sodium  hydroxide  can 
be  considered  as  water  in 
which  one  atom  of  the  metal 
hydrogen  has  been  replaced 
by  one  of  the  metal  sodium, 
and  similar  reactions  may  be 
expected  when  metals  are 
brought  in  contact  with  a 
compound  of  hydrogen  and 
any  negative  element  or 
group  of  elements'. 

Such  changes  result  with 
other  elements,  the  metallic  properties  of  which  are  as  pronounced 
as  those  of  sodium ;  for  instance,  potassium  will  react  with  water 
as  follows  : 10  — 

K  +HOH  =  KOH  +  H. 

Potassium  +  water    =  Potassium  hydroxide  +  hydrogen. 

The  group  of  elements  OH,  when  it  is  united  with  some  element  or 
group  of  elements,  is  called  the  hydroxyl  group,  and  substances 
which  contain  this  are  called  hydroxides ;' thus,  water, 'HOH,  could 
be  called  hydrogen  hydroxide,  and  by  replacing  the  hydrogen  with 
some  other  element,  such"  as  sodium,  potassium,  calcium,  or  mag- 
nesium, we  have  sodium,  potassium,  calcium,  or  magnesium  hy- 
droxides. 

NaOH,  Sodium  hydroxide, 

K  OH,  Potassium  hydroxide, 
Ca  (OH  )2 ,  Calcium  hydroxide. 
Mg(OH)2,  Magnesium  hydroxide. 
Fe  (OH  )2 ,  Ferrous  hydroxide. 
Zn  (OH  )2 ,  Zinc  hydroxide. 

These  hydroxides  vary  in  formula,  by  reason  of  the  fact  that  the  in- 
dividual atoms  of  different  metals  are  capable  of  uniting  with  differ- 


30  HYDROXIDES. 

ent  numbers  of  hydroxyl  groups.  In  the  examples  before  us  we 
have  two  elements,  sodium  and  potassium,  the  atoms  of  which  unite 
each  with  one  hydroxyl  group ;  two  elements,  calcium  and  magne- 
sium, where  they  unite  with  two ;  and  we  can  cite  a  third  class 
of  elements,  of  which  iron  and  aluminium  are  representatives,  in 
which  one  atom  of  the  metal  can  unite  with  three  of  these  groups. 

Al  (OH) 3 ,  Aluminium  hydroxide, 
Fe  (OH)8,  Ferric  hydroxide. 

The  hydroxides  of  the  metals,  as  well  as  the  oxides,  are  bases, 
and  with  acids  they  yield  salts  and  water;  they  therefore  present 
the  salne  chemical  contrast  toward  anhydrides  of  acids,  or  acids,  as 
do  the  oxides  (see  page  25),  thus : 

K20  +S03  =  K2S04 

Potassium  oxide  +  Sulphuric  anhydride  =  Potassium  sulphate. 

2KOH  +S03  =K2S04  +H20 

Potassium  hydroxide  -f  Sulphuric  anhydride  =  Potassium  sulphate  +  Water. 

K2O  +  H2S04  =K2S04  '  +H20 

Potassium  oxide  +  Sulphuric  acid  =  Potassium  sulphate  -}-  "Water. 

2KOH  +  H2S04  =K2S04  +2H20 

Potassium  hydroxide  +  Sulphuric  acid  =  Potassium  sulphate  +  Water. 

The  acid  employed  in  the  last  two  reactions  is  formed  from  sul- 
phuric anhydride  and  water,  S03  +  H2  0  =  H2  S04 ;  so  that  both 
the  oxide  and  hydroxide  of  potassium,  with  either  the  anhydride  of 
sulphuric  acid  or  the  acid  itself,  will  yield  a  salt  (potassium  sul- 
phate) and  water  (excepting  in  the  case  of  the  oxide  of  potassium 
and  sulphuric  anhydride  when  no  water  is  formed).  What  is  true 
of  potassium  is  true  of  other  metals  as  well.  The  oxides  and 
hydroxides  of  the  metals  are  therefore  bases,  if,  with  acids,  they 
yield  salts  and  water.  The  hydroxyl  group  is  not  confined  in  its 
union  to  metals  alone,  the  not-metals  likewise  are  capable  of  form- 
ing hydroxides,  but  the  hydroxides  of  not-metals  are  acids  ;  the 
discussion  of  the  character  of  these  hydroxides  is,  however,  better 
deferred  until  the  pupil  becomes  acquainted  with  a  greater  number 
of  chemical  compounds. 

The  more  pronouncedly  metallic  in  its  nature  a  metal  is,  the 
more  readily  will  it  be  able  to  decompose  water,  liberating  hydro- 
gen. Sodium,  potassium,  or  calcium  does  so  at  ordinary  tempera- 
tures ;  magnesium  does  so,  provided  the  water  is  boiling ;  iron,  if 


HYDROGEN  ;    PREPARATION.  31 

the  metal  is  heated,  for  when  steam  is  passed  over  red  hot  iron, 
hydrogen  is  produced,  as  follows  : 

3  Fe  +  4  H2  0  =  Fe3  04  +  8  H. 

Not-metals  (excepting  red  hot  carbon)  do  not  generate  hydrogen 
from  water. 

3.  By  the  decomposition  of  acids  by  metals. 

A  better  method  of  preparing  hydrogen  for  laboratory  use  than 
the  one  by  the  actions  of  metals  on  water  is  by  that  of  the  metals 
011  acids,  by  which  means  hydrogen  and  a  salt  are  formed,  thus  :  — 


As  action  of  sodium  on  hydrochloric  acid  is  not  practically 
available,  it  is  necessary  to  substitute  some  other  metal,  such  as 
zinc  or  iron,  when  the  chloride  of  the  metal  and  hydrogen  are 
formed  :  — 

Zn   +  2HC1  =ZnCl2  +  2H 

Zinc  +  Hydrochloric  acid  =  Zinc  chloride        +  Hydrogen.f 

Fe    +  2HC1  =FeCl2  +  2H 

Iron  +  Hydrochloric  acid  =  Ferrous  chloride  +  Hydrogen. 
Instead  of  hydrochloric  acid,  dilute  sulphuric  acid  could  be  used  ; 

tlms:       Zn  +H2S04  =ZnS04  +2H 

Zinc  +  Sulphuric   acid  =  Zinc  sulphate         -{-  Hydrogen. 

Fe    +H2S04  =FeS04  +  2H 

Iron  +  Sulphuric  acid  =  Ferrous  sulphate  +  Hydrogen. 

The  action  of  either  sodium  or  potassium  on  water  does  not 
differ  in  principle  from  that  of  zinc  or  iron  on  hydrochloric  or  sulv 
phuric  acid,  for  in  the  case  of  either  water  or  the  acids,  the  me- 
tallic element,  hydrogen,  is  attached  to  a  negative  element,  or  group 
of  elements,  thus  :  — 

H(OH),  H2(S04),  H(C1), 

*  This  reaction  can  take  place  between  sodium  and  gaseous  H  Cl.  (See 
note  32  appendix.  )  The  solution  of  hydrochloric  acid,  which  is  the  ordinary 
laboratory  preparation,  is  not  available  for  the  purpose,  for  the  action  is  too 
violent. 

t  On  comparing  the  formulae  of  zinc  chloride  and  ferrous  chloride  with 
those  of  zinc  hydroxide  and  ferrous  hydroxide,  the  pupil  will  see  that  one 
atom  of  the  metal  combines  with  the  same  number  of  chlorine  atoms  as  it 
does  with  hydroxyl  groups. 


32 


HYDROGEN  ;   PREPARATION. 


and  this  hydrogen,  when  either  the  hydroxide  or  a  salt  is  formed, 
is  displaced  by  another  metal  which  possessed  greater  chemical 
energy  when  in  contact  with  these  negative  groups.  This  is  true 
of  sodium  and  hydroxyl,  zinc  and  the  group  S04,  or  zinc  and 
chlorine  ;  for  in  each  case  the  heat  of  formation  of  the  hydroxide, 
sulphate  or  chloride,  is  greater  than  is  that  of  the  corresponding 
hydrogen  compound  ;  as  a  consequence,  the  metal,  when  brought  in 
contact  with  those  compounds,  will  tend  to  displace  the  hydrogen 
to  form  a  system  possessing  less  energy  than  did  the  original  one  ;  * 
thus  in  the  reaction, 


the  Zn  -f-  H2  S04  possesses  more  chemical  energy  than  Zn  S04  -+-  2  H, 

as  is  seen  by  the  fact  that  heat  is  pro- 
duced when  the  reaction  takes  place. 

In  order  to  prepare  hydrogen  for  labo- 
ratory use,  it  is  best  to  place  granulated 
zinc  n  in  a  flask,  with  delivery  and  safety 
tube   (Fig.  5),  and  to   pour  in 
dilute  sulphuric  12  acid  through 
the  thistle  tube.     The  gas  gen- 
erated from  zinc  is  purer  than 
that  which   is  prepared    from 
ordinary  iron,  so  that  the  latter 
is  not  used  for  laboratory 
purposes.13'  u 

Hydrogen  is  a  color- 
less gas  ;  when  pure  it  is 
odorless  and  tasteless  ;  at 
probable  temperature 
of  —  213°  and  a  pressure 
of  40  atmospheres,  it  is 
changed  to  a  colorless 
liquid.  It  is  not  very 


Fig.  5. 


soluble  in  water,  one  hundred  volumes  of  that  substance  dissolv- 

*  In  discussing  chemical  energy,  if  the  reaction  under  consideration  is  ex- 
pressed in  the  form  of  a  chemical  equation,  the  bodies  to  the  left  of  the  sign  of 
equality  are  termed  "  a  system,"  and  those  to  the  right  are  given  the  same 
name.  Pattison  Muir,  Principles  of  Chemistry,  p.  247. 


HYDROGEN  ;   PROPERTIES.  33 

ing  1.82  volumes  of  hydrogen  at  20°.*  Hydrogen  has  the  smallest 
specific  gravity  of  any  substance  with  which  we  are  acquainted ;  it 
occupies  14£  times  the  space  taken  by  an  equal  weight  of  air,  or 
11,160  times  that  of  water,  its  specific  gravity  being  .06960.  In 
consequence  of  this  specific  lightness,  jars  can  be  filled  with  hydro- 
gen by  placing  them  mouth  downward,  and  allowing  the  gas  to 
enter  through  a  tube  extending  upward  to  the  bottom,  when  tjie 
hydrogen  will  expel  the  air ;  the  collecting  of  gases  in  this  manner 
is  termed  "  upward  displacement."  Balloons  or  soap  bubbles  filled 
with  hydrogen  will  rise  in  the  air,  and  if  two  beakers  are  suspended, 
mouth  downward,  on  a  balance,  and  then  exactly  tared,  one  arm  of 
the  balance  can  be  caused  to  rise  by  filling  the  corresponding 
beaker  with  hydrogen,  poured  upward  from  another  vessel. 

Hydrogen  passes  readily  through  porous  substances.  It  has  the 
smallest  specific  gravity,  and  the  greatest  rate  of  diffusion  of  any 
gas,  because  the  rate  of  diffusion  of  gases  is  inversely  as  the  square 
roots  of  their  densities, f  the  law  being  approximately  true.  By 
diffusion  we  mean  the  power  which  a  gas  has  of  completely  mixing 
with  another  gas,  or  gases,  which,  without  stirring,  are  placed  in 
contact  with  it.  This  can  take  place  even  though  the  gases  are 
separated  by  a  partition,  provided  this  latter  is  sufficiently  porous 
to  allow  of  the  passage  of  the  molecules.  Hydrogen  can  pass 
through  substances  such  as  unglazed  earthenware,  paper,  or  even 
metals  like  platinum,  when  the  latter  are  heated.15  Some  metals,  for 
example,  palladium,  have  the  power  of  condensing  large  quantities 
of  hydrogen.  Palladium,  when  in  the  shape  of  wire,  can  absorb 
872  times  its  own  volume,  and,  when  used  as  a  negative  electrode 
in  a  battery,  it  can  condense  very  nearly  1,000  volumes ;  so  that  1 
c.c.  of  palladium  can  absorb  about  a  liter  of  the  gas.$  The  palla- 
dium gains  in  volume  during  the  operation,  and  becomes  specifically 
lighter.  Hydrogen  so  absorbed  is  said  to  be  occluded.  Occluded 
hydrogen  is  chemically  much  more  active  than  is  the  ordinary  gas, 
and  reacts  with  substances  on  which  the  element  ordinarily  has  no 
effect.16  Heat  is  evolved  in  the  process  of  occlusion ;  and  a  portion 
of  the  hydrogen  taken  up  by  palladium  is  probably  chemically  com- 

*  Cl.  Winckler,  Berichte  d.  Deutsch.  Chem.  Gesell.  24;  97. 
t  The  terms  " density  "  and  "specific  gravity,"  when  referring  to  gases, 
are  used  interchangeably. 

J  M.  Thoma,  Zeitschrift  fur  Physikalische  Chemie ;  3,  83. 


34  HYDROGEN  ;   DIFFUSION,    OCCLUSION. 

bined ;  the  metal  liberates  a  portion  of  its  occluded  hydrogen  even 
at  ordinary  temperature,  while  at  130°  to  140°  the  greater  part 
passes  off.  So  energetically  does  the  hydrogen  occluded  by  palla- 
dium act  that,  if  the  metal  is  saturated  with  the  gas,  it  will  glow 
spontaneously  in  the  air,  the  hydrogen  uniting  with  oxygen  to 
form  water.  The  condensation  of  gases  by  metals  can  be  easily 
shown  by  suspending  a  coiled  platinum  wire  in  a  burning,  not- 
luminous  gas  flame  until  it  is  red  hot,  turning  out  the  flame  and 
then  instantly  turning  it  on  again.  The  platinum  occludes  the 
gases  escaping  from  the  burner,  and,  being  warm,  they  unite  with 
the  oxygen  of  the  atmosphere,  so  that  the  wire  continues  to  glow 
and  finally  kindles  the  gas.  Occluded  gases  are  frequently  used  in 
chemical  operations. 

Animals,  when  placed  in  an  atmosphere  of  hydrogen,  die  of 
asphyxia ;  hydrogen  is  not  poisonous,  per  se,  as  is  proved  by  the 
fact  that  animals  can  live  without  discomfort  in  an  atmosphere  of 
hydrogen  to  which  has  been  added  a  sufficient  supply  of  oxygen. 
Hydrogen  is  not  able  to  support  combustion  in  the  ordinary  sense 
of  the  term,  a  burning  candle  being  extinguished  when  placed  in  it. 
In  order  to  be  combustible  in  a  gas,  a  body  must  have  a  great  ten- 
dency to  unite  with  that  gas ;  necessarily  the  substances  which  are 
ordinarily  considered  as  combustible,  such  as  wood  or  a  candle,  can- 
not burn  in  hydrogen,  for  they  themselves  are  largely  composed  of 
hydrogen.  Hydrogen  burns  quite  readily  in  oxygen,  with  a  very 
hot  flame,  and  vice  versa,  oxygen  will  burn  in  hydrogen  with  the 
production  of  an  equally  hot  flame ;  in  both  instances  water  is  pro- 
duced (see  page  36).  Hydrogen  will  also  burn  in  other  gases  for 
which  it  has  a  great  affinity,  in  chlorine,  for  instance.  The  hydrogen 
compounds  of  the  elements  in  which  the  gas  is  burned  will,  in  each 
instance,  be  produced. 

In  order  to  prove  the  formation  of  water  by  the  combustion  of 
hydrogen  in  the  air,  we  have  but  to  dry  thoroughly  the  gas  passing 
from  a  generator ; 17  ignite  it  at  the  tip  of  a  burner, 18  and  then  hold 
a  cold  beaker  over  the  same.  Drops  of  water  soon  collect  and  run 
down  the  sides  of  the  vessel.  The  flame  of  hydrogen  burning  in 
oxygen  has  a  very  high  temperature,  and  is  very  nearly  not-lumi- 
nous ;  but  it  can  be  changed  to  a  luminous  one  by  dropping  a  little 
powdered  coal  or  lycopodium  into  it,  for  not-luminous  flames  are 
simply  glowing  gases :  they  ara  made  luminous  by  glowing  solids 


HYDROGEN;  COMBUSTION  OF.  35 

within  them.  Hydrogen,  like  other  substances,  burns  more  energet- 
ically in  oxygen  than  in  air,  so  that  the  temperature  of  a  flame  of 
hydrogen  burning  in  the  former  is  much  higher  than  that  of  the 
same  gas  burning  in  the  latter.  This  fact  is  taken  advantage  of  in 
the  use  of  the  oxy-hydrogen  blow-pipe.  This  apparatus  consists  of 
two  concentric  tubes  drawn  nearly  to  a  point ;  through  the  outer  one 
hydrogen  is  admitted  and  ignited,  and  then  a  stream  of  oxygen  gas 
is  passed  through  the  central  tube,  thus  mixing  the  two  gases  at  their 
point  of  combustion.  The  most  extreme  heat  of  a  flame  of  hydrogen 
burning  in  oxygen  is  thus  obtained.  The  illuminating  power  of  cal- 
cium lights,  is  due- to  a  piece  of  lime  (an  infusible  substance),  heated 
to  incandescence  by  the  oxyhydrogen  blow  pipe ;  zinc,  iron,  or  tin, 
burns  readily  in  the  flame ;  a  piece  of  platinum  wire  is  instantly 
melted;  indeed,  the  fusing  of  platinum  for  the  manufacture  of 
various  utensils  is  accomplished  by  this  means. 


36  WATER;  CHEMICAL  HISTORY. 


CHAPTER   IV. 

WATER. 

Symbol,  H20;  specific  gravity  1,  at  4°  C ;  specific  gravity  of  vapor, 
air  =  1,  is  .6208  ;  H2  =  2,  is  17.88. 

WHEN  hydrogen  is  burned  in  oxygen,  or  oxygen  in  hydrogen, 
water  is  formed.  The  study  of  the  composition  of  this  substance, 
and  the  methods  by  which  this  has  been  advanced,  might  serve  as  a 
type  of  all  other  similar  investigations ;  so  that  for  the  purpose  of 
becoming  somewhat  acquainted  with  the  means  of  chemical  research, 
the  subject  will  be  discussed  at  some  length.  That  hydrogen  and 
oxygen,  in  uniting,  form  water  is  easily*proved ;  but  the  question  at 
once  arises,  is  water  alone  produced,  or  may  not  some  other  sub- 
stance originate  simultaneously,  when  the  combustion  takes  place  ? 
Indeed,  when  all  of  the  products  of  combustion  of  hydrogen  in  the 
air  are  collected,  an  acid  substance,  nitric  acid,  is  also  found  to  be 
present.  The  formation  of  an  acid  by  this  means  puzzled  the 
original  investigators ;  and  not  until  they  had  exploded  a  mixture 
of  pure  hydrogen  and  pure  oxygen,  as  Cavendish  did  in  1780,  and 
so  discovered  that  no  acid  whatever  was  produced  under  those  cir- 
cumstances, did  they  come  to  the  conclusion  that  hydrogen  and 
oxygen,  in  burning,  formed  water  only.  Another  problem  as 
regards  water  then  presented  itself.  Before  the  last  quarter  of 
the  eighteenth  century,  it  was  generally  supposed  that  water,  by 
boiling,  was  changed  to  an  earthy  substance,  for  on  boiling  even 
pure  distilled  water  in  glass  or  earthen  vessels,  an  earthy  residue 
remained  after  evaporation.  Lavoisier  took  upon  himself  to  prove 
that  the  universally  accepted  notion  was  erroneous.  In  order  to  do 
this,  he  sealed  some  pure  water  in  a  glass  flask,  weighed  it,  then 
kept  it  heated  for  some  weeks  and  found  that  the  total  weight  was 
unchanged.  On  evaporating  the  water,  a  solid  substance  remained, 
but  then  Lavoisier  found  that  the  flask  had  lost  in  weight,  and  this 
loss  in  weight  was  exactly  equal  to  that  of  the  earthy  residue.  As 
a  consequence,  he  came  to  the  conclusion  that  the  earthy  substance, 


WATER  ;    COMPOSITION   OF.  37 

supposed  to  be  formed  by  the  boiling  of  water,  was  nothing  more 
than  a  dissolved  portion  of  the  glass,  deposited  by  evaporation. 
Scheele,  at  the  same  time,  showed  that  this  earthy  residue  had 
exactly  the  same  composition  as  the  glass  in  which  the  water  had 
been  boiled.  Chemists  had  then  proved  two  things  by  their  inves- 
tigations ;  that  hydrogen  and  oxygen,  in  uniting,  form  nothing  but 
water,  and  that  this  water,  in  evaporating,  was  volatilized  unchanged. 
We  have  seen  that  water  is  decomposed  by  the  electric  current. 
It  is  not  only  necessary  for  the  chemist  to  prove  that  hydrogen  and 
oxygen  are  produced  by  this  means  ;  he  must  also  show  that  nothing 
but  these  gases  is  formed.  Sir  Humphry  Davy  noticed  that,  in 
electrolyzing  water,  a  certain  amount  of  alkali  separated  at  the 
negative  pole,  while  acid  was  formed  at  the  positive  one.  There 
were  only  three  possible  theories  to  hold :  one  was  that  these 
substances  were  produced  from  the  water;  the  second,  that  they 
were  formed  by  the  decomposition  of  the  glass,  for  glass  is  composed 
of  an  acid  and  an  alkaline  substance ;  and  the  third,  that  the  sur- 
rounding air  took  part  in  the  change.  In  order  to  discover  the 
source  of  these  impurities,  Davy  transferred  the  water  to  a  gold 
vessel,  and  then  continued  the  process  of  electrolysis.  Acid  and 
alkali  were  formed  as  before,  but,  while  the  production  of  acid  con- 
tinued, the  amount  of  alkali  became  less  and  less,  until  it  finally  dis- 
appeared, showing  that  it  was  owing  to  the  decomposition  of  the 
glass.  Davy  came  to  the  conclusion  that  this  latter  was  due  to 
nitric  acid,  produced  from  the  atmosphere,  and,  as  a  consequence, 
the  gold  vessel  was  placed  under  a  bell-jar  on  the  receiver  of  an  air 
pump,  the  air  exhausted  and  replaced  by  pure  hydrogen,  when  the 
formation  of  acid  ceased.*  It  was  proved,  therefore,  that  water 
changed  to  nothing  but  hydrogen  and  oxygen  on  electrolysis,  and 
with  that  the  chain  of  evidence  was  complete.  There  remained 
now  to  determine  what  proportionate  volumes  of  hydrogen  and 
oxygen  were  produced  in  the  electrolysis  of  water,  and  what  pro- 
portionate volumes  united. to  form  water.  This  matter  had  been  111 
dispute  for  a  number  of  years,  until  Gay  Lussac  and  Humboldt 
showed  that  exactly  twice  as  much  hydrogen  as  oxygen,  by  volume, 
is  contained  in  water ; 19  in  other  words,  for  every  two  cubic  centi- 
meters of  hydrogen  there  is  one  cubic  centimeter  of  oxygen,!  and 

*  See  Kopp,  Geschichte  der  Chemie;  1,  377. 

t  See  E.  W.  Morley,  Amer.  Journal  of  Science,  [3]  41;  220,  276. 


38  WATER;  COMPOSITION  BY  WEIGHT. 

if  hydrogen  and  oxygen  are  mixed  in  exactly  these  proportions,20 
the  gas  produced  explodes  violently  when  ignited  by  means  of 
a  taper  or  electric  spark.  If  the  exact  mixture  is  exploded  in  a 
short  heavy  glass  tube  closed  with  mercury,  over  a  trough  filled 
with  that  metal,  then,  after  the  explosion,  the  mercury  will  com- 
pletely fill  the  tube,  as  the  volume  of  water  formed  is  insignificant 
as  compared  to  the  volume  of  the  gases  used.  The  following  facts 
have  then  been  proved  as  regards  water.  It  is  composed  of  hydro- 
gen and  oxygen  only;  it  decomposes  into  nothing  but  hydrogen 
and  oxygen ;  the  volume  of  hydrogen  is  to  the  volume  of  oxygen  as 
2:1;  and  furthermore,  two  volumes  of  hydrogen  mixed  with  one  of 
oxygen  form  an  explosive  mixture,  from  which  only  water  is 
produced,  without  a  residue  of  either  hydrogen  or  oxygen. 

Having  proved  the  composition  of  water  by  volume,  our  next 
task  is  to  discover  the  same  by  weight,  and  this  was  first  most 
accurately  determined  by  Dumas.*  By  passing  pure  hydrogen 
over  heated  copper  oxide,  the  following  reaction  takes  place :  — 

CuO  +2H  =  Cu          +H20 

Copper  oxide       +  hydrogen      =  Copper     +  water. 

By  this  means  we  have  a  most  ready  method  of  discovering  the 
proportional  parts  by  weight  in  which  hydrogen  and  oxygen  unite. 
For,  let  us  suppose  the  copper  oxide  to  be  perfectly  dry  and 
accurately  weighed,  then,  after  the  reaction,  the  weight  of  the 
remaining  copper  subtracted  from  that  of  the  copper  oxide  would 
give  us  the  weight  of  oxygen  which  has  gone  to  form  water ;  if,  by 
some  means,  we  can  collect  the  water  formed,  and  weigh  the  same 
with  equal  accuracy,  the  difference  between  the  weight  of  the 
oxygen  and  that  of  the  total  water  would  give  the  weight  of  hy- 
drogen. 

a  =  weight  of  copper  oxide.  c  =  weight  of  water  formed, 

b  =  weight  of  copper  after  the  re-      c  —  x  =  weight  of  hydrogen. 

action, 
a  —  b  =  weight  of  oxygen  =  x  and  x  :  c  —  x  will  represent  the  ratio 

in  which  oxygen  and  hydrogen  unite  to  form  water. 

This  ratio  was  found  by  Dumas  to  be  1  :  7.98.  Dumas  took 
great  care  to  purify  the  hydrogen  which  he  used.21  The  water  was 

*  Berzelius  and  Dulong  first  undertook  to  determine  the  quantitative  com- 
position of  water  by  this  method. 


WATER  ;    FORMULA   OF.  39 

collected  by  passing  the  gas  through  an  empty  glass  flask,  in  which 
the  greater  part  was  condensed,  and  then  through  a  series  of  tubes 
filled  with  fused  caustic  potash  and  phosphorus  pentoxide,  which 
substances  absorbed  all  of  the  remainder.  The  tubes  were  filled 
with  perfectly  dry  and  pure  air  before  the  operation,  and  then 
weighed ;  after  the  water  had  been  formed  they  were  once  more 
filled  with  dry  and  pure  air  and  again  weighed,  when  the  gain 
in  weight  gave  the  amount  of  water  formed.  This  latter,  in  nine- 
teen experiments,  was  945.439  grams;  the  loss  in  weight  of  the 
copper  oxide,  and  consequently  the  amount  of  oxygen  in  the  water, 
was  840.161  grams,  leaving  for  hydrogen  (945.439  —  840.161)  or 
105.278  grams.  The  ratio  between  the  combining  weights  of  hydro- 
gen and  oxygen,  therefore,  is  as  105.278  :  840.161,  or  as  1  :  7.98.  Of 
late  Morley,  by  using  a  different  method,  has  proved  that  the  ratio 
is  more  probably  1  :  7.94.* 

The  facts  which  we  have  discovered  as  regards  water  are  then 
as  follows : 

It  can  be  decomposed  into  hydrogen  and  oxygen,  and  yields  two 
volumes  of  hydrogen  to  one  of  oxygen ;  it  is  formed  from  two  vol- 
umes of  hydrogen  and  one  of  oxygen,  and  when  the  water  so  formed 
is  heated  above  the  boiling  point  (i.e.,  measured  as  steam),  it  will  be 
seen  that  2  volumes  of  hydrogen  -f  1  volume  oxygen  produce  2  vol- 
umes water  vapor ;  and,  furthermore,  water  contains  the  two  ele- 
ments in  the  proportional  parts  by  weight  of  one  of  hydrogen  to 
eight  of  oxygen.  AVe  shall  subsequently  see  that  in  equal  volumes 
of  gases,  under  the  same  conditions  of  temperature  and  pressure, 
there  are  equal  numbers  of  particles ;  so  that  if  two  volumes  of 
hydrogen  unite  with  one  of  oxygen  to  form  water,  we  are  tolerably 
safe  in  assuming  that  two  atoms  of  hydrogen  unite  with  one  of 
oxygen,  and  if  we  neglect  the  decimals,  two  parts  by  weight  of 
hydrogen  unite  with  sixteen  of  oxygen ;  from  this  it  follows  that  in 
all  probability  the  atomic  weight  of  oxygen  is  sixteen.  This  last 
conclusion  cannot  be  absolutely  certain,  as  we  shall  subsequently 
see ;  but  we  are  sure  that  one  part  by  weight  of  hydrogen  unites 
with  eight  of  oxygen  to  form  water,  and  formerly  chemists  were  of 
the  opinion  that  speculation  should  not  extend  beyond  this  certainty. 
They  wrote  the  formula  of  water  HO,  meaning  by  this  that  hydro- 

*  Very  nearly  the  same  ratio  (7.935  to  7.945)  has  been  found  by  Kayleigh, 
Keiser,  Noyes,  Leduc,  and  others. 


40  WATER;  PROPERTIES. 

gen  and  oxygen  unite  to  form  water,  in  the  proportion  of  one  part 
by  weight  of  hydrogen  to  eight  of  oxygen ;  and  as  eight  parts  by 
weight  of  oxygen  are  equivalent  in  combining  power  with  one  of 
hydrogen,  eight  was  called  the  equivalent  weight  of  oxygen.  The 
custom  in  respect  to  oxygen  was  followed  in  regard  to  other  elements 
as  well.  In  despair  of  finding  some  guide  by  which  to  determine 
the  true  atomic  weights,  equivalent  weights  were  resorted  to,  as  the 
latter  involved  no  speculation,  but  were  simply  founded  upon 
the  known  experimental  facts.  As  chemical  and  physical  theory 
and  investigation  became  more  perfect,  it  was,  however,  seen  that 
we  could  form  very  definite  conclusions  as  regards  the  atomic 
weights,  so  that  equivalent  weights  were  abandoned,  and  our 
present  atomic  weights  came  into  general  use.  The  investigations 
on  the  composition  of  water  will  serve  to  illustrate  the  means 
by  which  our  present  chemical  facts  have  become  known,  and  the 
pupil  should  remember  that  when  in  the  future  the  atomic  weights 
of  elements  or  the  formulae  of  compounds  are  mentioned,  these  have 
been  discovered  by  some  equally  painstaking  means  of  investigation. 
Water,  at  ordinary  temperatures,  is  a  nearly  colorless  liquid ; 
when  light  is  passed  through  a  thick  layer  of  the  substance,  it 
will  be  seen  to  have  a  distinctly  blue  color.  When  heated  to  its 
boiling  point,  which,  with  760  m-m-  pressure,  is  at  100°  centigrade, 
it  is  transformed  into  a  colorless  gas..  When  water  is  cooled,  the 
substance  contracts,  following  the  general  law ;  when  heated,  it 
expands.  If,  however,  water  having  a  temperature  above  4°  centi- 
grade is  cooled,  it  will  contract  until  that  temperature  is  reached, 
and  will  then  begin  to  expand  until  0°  centigrade,  at  which  point  it 
freezes.  Water  at  4°  centigrade,  then,  has  the  greatest  specific  grav- 
ity,* so  that  any  body  of  water,  on  cooling  toward  that  temperature, 
will  become  specifically  heavier ;  that  portion  on  the  surface,  because 
it  cools  firstr  will  sink ;  the  warmer  and  lighter  water  below  will 
rise,  and  in  this  manner  a  continuous  circulation  will  be  kept  up 
until  the  entire  body  has  arrived  at  the  temperature  of  4°  centi- 
grade. Now,  the  water  on  cooling  further  will  expand,  and  hence 
the  cooler  portion  will  float  on  the  surface  until  it  arrives  at  0°, 
when  freezing  begins,  and  the  crust  of  ice  formed,  not  being  able 
to  sink,  will  act  as  a  protection  to  the  water  below.  As  a  con- 

*  If  we  place  the  density  of  water  at  4°  centigrade  at  1.,  then  water  at  0° 
centigrade  has  a  specific  gravity  of  .99988. 


WATER  ;    SOLUTION.  41 

sequence,  water  freezes  on  the  surface  and  not  from  below  upward. 
In  freezing,  water  expands,  the  specific  gravity  of  ice  being  .9167  ; 
on  cooling  below  its  freezing  point,  ice  continuously  contracts,  fol- 
lowing the  usual  law.  The  freezing  point  of  water,  or,  rather,  the 
melting  point  of  ice,  at  standard  atmospheric  pressure,  is  taken  as 
0°  on  the  centigrade  thermometer ;  this  point  is  lowered  by  com- 
pression, so  that  ice  at  0°  centigrade  can  be  fused  by  increasing  the 
pressure ;  as  soon  as  this  is  relieved  the  liquid  instantly  freezes.  This 
fact  can  be  experimentally  proved  by  firmly  pressing  together  two 
pieces  of  ice,  which  will  adhere  as  soon  as  the  pressure  is  relieved. 

Water  is  able  to  dissolve  a  large  number  of  substances  to  form 
solutions.  Solutions  are  homogeneous  mixtures  of  two  or  more 
substances,  one  of  which  must  be  either  gaseous  or  liquid.  These 
mixtures  cannot  be  separated  by  simple  mechanical  means.  Gases 
can  form  these  homogeneous  mixtures  in  any  proportion.  When  two 
substances  are  liquid,  the  solution  may  take  place  in  any  proportion, 
as  between  alcohol  and  water,  alcohol  and  ether,  acetic  acid  and 
water  ;  or  one  liquid  may  partially  dissolve  another,  as  in  the  case 
of  water  and  ether ;  *  or,  lastly,  one  liquid  may  be  entirely  insoluble 
in  another,  as  in  the  case  of  some  oils  and  water.  When  a  solid 
dissolves  in  a  liquid,  the  latter  is  only  able  to  take  up  a  certain 
quantity,  which,  with  any  given  liquid,  varies  with  the  nature  of 
the  solid.  When  the  liquid  has  dissolved  as  much  of  the  solid  as 
it  will,  the  solution  is  said  to  be  saturated.  As  a  rule,  gases  are 
less  soluble  in  liquids  the  higher  the  temperature ;  they  are  there- 
fore expelled  from  their  solutions  on  heating.  Solids,  on  the  other 
hand,  are,  as  a  rule,  more  soluble  in  hot  liquids  than  in  cold  ones. 
As  a  consequence,  a  saturated  solution,  at  the  boiling  point  of  the 
liquid,  will  contain  more  of  the  dissolved  solid  than  it  will  at 
a  lower  temperature,  so  that,  on  cooling,  the  dissolved  substance 
separates,  frequently  in  a  crystalline  form.  The  process  of  dissolv- 
ing crystalline  solids  in  hot  liquids  and  separating  by  cooling,  is 
called  recrystallization,  and  is  very  frequently  employed  as  a 
means  of  purifying  crystalline  substances.  The  solubility  of  solids 

*  In  this  latter  case  we  can  scarcely  say  that  one  liquid  dissolves  the  other, 
for  to  take  an  example  from  the  instance  just  cited,  the  ether  will  dissolve  just 
as  much  water  as  the  water  will  dissolve  ether.  For  a  more  extended  element- 
ary discussion  of  solutions  see  Ostwald's  Outlines  of  General  Chemistry,  Walker's 
translation;  Macmillan,  1890. 


42  WATER   OF   CRYSTALLIZATION. 

in  water  varies  very  greatly ;  some  substances  are  insoluble,  others 
but  very  slightly  so,  so  that  the  solution  may  be  saturated  when 
only  a  trace  of  dissolved  substance  is  present ;  again,  we  have  solids 
such  as  potassium  hydroxide,  calcium  chloride,  or  magnesium  chlo- 
ride, which  are  soluble  even  in  their  own  volume  of  water.  The 
solutions  formed  are  clear  liquids,  which  may  have  the  color  of  the 
dissolved  solid.  In  making  solutions  we  frequently  have  marked 
changes  of  temperature.  Where  substances  form  no  chemical 
union  with  water  in  dissolving,  the  temperature  is  lowered,  for 
work  must  be  done  to  change  the  crystallized  solid  to  a  liquid ; 
thus  potassium  or  ammonium  nitrate,  on  dissolving  in  water,  effect 
a  very  marked  cooling. 

Quite  a  number  of  substances  can  chemically  take  up  water  to 
form  crystalline  compounds  of  a  definite  form ;  this  is  done  in  such 
a  way  that  a  given  quantity  of  the  solid  unites  with  a  definite  quan- 
tity of  water.  The  water  is  apparently  not  present  as  such,  for  the 
crystals  are  perfectly  dry,  the  water  appearing  only  when  the  solid 
is  heated,  for  it  then  passes  off.  Such  combined  water  is  known  as 
water  of  crystallization ;  the  number  of  molecules  of  water  of  crys- 
tallization united  with  one  formula  weight  of  the  solid,  as  well  as 
the  crystalline  form  of  the  substance,  are  always  the  same  for  any 
given  compound ;  in  some  few  cases,  however,  the  same  individual 
can  crystallize  in  different  forms,  with  different  amounts  of  water 
of  crystallization.  To  cite  a  few  examples,  one  formula  weight  of 
copper  sulphate  crystallizes  with  five  molecules  of  water,  as 
Ou  S04 ,  5  H2  0  ;  ferrous  sulphate  with  seven  molecules  as  Fe  S04 , 
7  H2  0  ;  sodium  carbonate  with  ten  molecules,  as  Na2  C03 , 10  H2  0  ; 
and  alum  with  twenty-four  molecules,  as  A12  K2  (S04)4  -f-  24  H2  0. 
Substances  which  are  chemically  similarly  constituted  will  fre- 
quently crystallize  in  the  same,  or  at  least  in  very  similar, 
crystalline  forms ;  such  substances  are  said  to  be  isomorphous.*  A 

*  The  student  should  consult  some  elementary  text-book  in  regard  to  the 
principles  of  crystallography.  Isomorphous  salts  are,  for  instance, 

MgSO4,7H2O.  CaCO3. 

Zn  SO4 ,  7  H2  O.  Sr  CO3 . 

Two  or  more  substances  are  not  truly  isomorphous,  unless  one  can  replace 
the  other  without  altering  the  crystalline  form,  thus,  MgSO4,  7H2O  and 
Zn  SO4 ,  7  H2  O,  both  crystallize  in  the  rhombic  system,  and  zinc  can  replace 
magnesium  in  crystals  of  Mg  SO4 ,  7  H2  O,  without  altering  the  crystalline  form. 


WATEE  ;   CHEMICAL   UNION   OF.  43 

number  of  compounds  lose  their  water  of  crystallization  on  stand- 
ing in  the  air,  their  crystalline  form  is  destroyed,  and  they  finally 
disintegrate  to  form  not-crystalline  powders;  such  substances  are 
efflorescent.  On  the  other  hand,  we  have  a  large  number  of  bodies 
which  are  capable  of  taking  up  water  from  the  atmosphere,  and 
dissolving  in  the  moisture  so  concentrated ;  these  are  deliquescent. 
By  far  the  greater  number  of  substances  are  neither  deliquescent 
nor  efflorescent  at  ordinary  temperatures,  but  all  substances  with 
water  of  crystallization,  give  the  liquid  up  at  a  comparatively  mod- 
erate heat.  When  water  of  crystallization  has  been  expelled  from 
a  substance,  the  body  is  said  to  be  anhydrous.  Anhydrous  salts 
dissolve  in  water  with  the  evolution  of  heat ;  many,  like  calcium 
chloride,  are  capable  of  absorbing  moisture  readily,  and  as  a  conse- 
quence are  most  useful  for  drying  gases.  It  must  be  borne  in 
mind,  however,  that  the  majority  of  substances  which  crystallize, 
do  not  contain  water  of  crystallization. 

Water  combines  with  many  substances  when  it  does  not  enter 
as  water  of  crystallization ;  for  example,  the  oxides  of  the  metals, 
when  soluble  in  water,  unite  with  that  substance  to  form  the 
hydroxides :  — 

K20  -fH20  =  2KOH 

Potassium  Oxide  +  Water  =  Potassium  hydroxide. 

Ca  0  +  H2  0  =  Ca  (OH)2 

Calcium  Oxide  +  Water  =  Calcium  hydroxide.* 

In  such  cases  the  water  entering  into  combination  with  the 
oxide  is  decomposed,  it  is  no  longer  present  as  water  but  as 
hydroxyl,  to  which  group  of  atoms  attention  has  already  been 
called  on  page  30.  The  hydroxides,  of  course,  differ  most  mark- 

*  Note  again  the  difference  between  the  formulas  K  OH  and  Ca  (OH)2 . 
Here,  we  see  one  atom  of  calcium  has  the  power  of  retaining  twice  as  many 
hydroxyl  groups  as  one  of  potassium  does ;  a  relationship  exactly  similar  to 
that  existing  between  potassium  and  magnesium  hydroxides,  and  similar 
to  the  relationship  in  the  formula  of  the  chlorides,  Ca  C12 ,  K  Cl.  If  we  were  to 
call  chlorine  X,  hydroxyl  Y,  the  relationship  would  become  apparent  in 
Ca  X2,  Ca  Y2;  KX,  KY.  These  chemical  changes  are  rendered  more  ap- 
parent if  the  pupil  will  take  the  trouble  to  write  out  the  formulae  of  reactions, 
atom  for  atom,  as  has  been  done  above. 


44  WATER;  PURIFICATION  OF. 

edly  in  stability ;  some  are  very  readily  decomposed  by  heat  into 
water  and  the  oxide,  while  others,  for  example,  potassium  hy- 
droxide, can  be  raised  to  a  high  red  heat  without  changing  to 
the  oxide.  As  a  rule,  the  less  pronouncedly  metallic  the  element 
forming  the  hydroxide  is,  the  more  readily  will  that  hydroxide  be 
decomposed  by  heat.  For  example,  calcium  hydroxide  is  decomposed 
as  follows  :  — 

/OHJJaO/H  Ca(QH)2  =Ca0  +H2Q 

/  /    TT  Calcium  hydroxide    =     Calcium  oxide  -f-   "Water. 

\OH  /OH 


Water  can  unite  with  the  anhydrides  of  acids  to  form  hy- 
droxides (called  acids),  as  readily  as  it  can  enter  into  combination 
with  the  oxides  of  metals.  Attention  will  subsequently  be  called 
to  this  class  of  reactions. 

Pure  water  is  very  difficult  to  obtain,  owing  to  the  great  capa- 
city which  the  liquid  has  for  dissolving  various  substances.  The 
impurities  may  be  of  two  kinds,  those  mechanically  suspended,  and 
those  dissolved ;  while  the  dissolved  impurities  may  be  classed  under 
three  heads,  gaseous,  liquid,  and  solid.  The  mechanically  suspended 
impurities  may  be  removed  by  nitration,  that  is,  by  passing  the  water 
through  some  porous  substance,  such  as  unsized  paper  (so-called 
filter  paper),  or  through  unglazed  porcelain.  Those  dissolved  must 
be  removed  by  distillation.*  The  gaseous  impurities  cannot  all  be 
eliminated  by  this  means,  for  the  oxygen  and  nitrogen  of  the  atmo- 
sphere, as  well  as  the  other  impurities  in  the  air,  are  soluble  in 
water,  and  will  therefore  be  found  in  the  distilled  water.  Liquid 
impurities  may  be  removed  by  repeated  distillation,  provided  their 
boiling  point  is  not  too  near  that  of  water ;  where  the  boiling  points 
of  two  liquids  are  within  10°  of  each  other,  a  complete  separation 
by  distillation  is  impossible.  The  solid  impurities  remain  behind 
in  the  vessel  from  which  water  has  been  distilled.  If  distilled 
water  is  to  be  free  from  gases,  it  must  be  boiled  for  some  time  in 
bottles,  and  then  hermetically  sealed  before  allowing  to  cool. 

All  naturally  occurring  waters  are  more  or  less  impure,  the 
purest  being  rain-water  and  melted  snow,  but  even  these  contain 

*  Distillation  is  simply  the  process  by  which  the  steam  from  boiling  water 
is  collected  in  suitable  vessels,  and  condensed  to  a  liquid.  For  a  description 
of  the  process  of  distillation,  any  larger  text-book  can  be  consulted. 


WATER;  NATURAL  WATERS.          45 

such  solid  and  gaseous  substances  as  they  can  collect  in  passing 
through  the  atmosphere.  Rain-water,  in  falling  on  the  soil,  takes 
up  such  soluble  substances  as  are  contained  therein,  and  in  so  doing 
is  changed  to  spring-water.  Of  course,  the  dissolved  impurities  of 
spring-water  vary  greatly  with  the  nature  of  the  soil  through  which 
the  water  has  passed.  When  the  amount  of  dissolved  impurity  is 
not  very  great  and  is  mainly  calcium  carbonate  or  sulphate,  the 
water  is  fresh  water ;  when  quantities  of  salts  are  dissolved  from 
the  soil,  the  water  becomes  a  mineral  water.  Some  springs  are 
heated  as  they  issue  from  the  earth ;  these  are  called  thermal 
springs.  The  most  frequent  salts  in  mineral  waters  are  sodium 
and  magnesium  sulphates,  sodium  carbonate  and  carbonate  of  iron, 
while  gaseous  constituents  like  carbon  dioxide  and  sulphuretted 
hydrogen  are  also  met  with.  Eiver  water  necessarily  must  con- 
tain the  dissolved  impurities  which  were  in  the  springs  from  which 
the  stream  has  its  source,  with  some  modifications  introduced  by  the 
nature  of  the  soil  over  which  it  has  passed ;  it  is  often  rendered 
unfit  for  drinking  purposes  by  contaminations  which  have  been  in- 
troduced, purposely  or  accidentally,  through  decaying  animal  or 
vegetable  substances.  As  streams  pass  through  more  or  less  thickly 
inhabited  regions,  the  proximity  of  dwellings  and  factories  is  apt  to 
cause  dangerous  impurities  in  the  water,  for  as  a  rule  very  little  care 
is  exerted  to  keep  sewage  from  reaching  the  same  stream  from  which 
drinking  water  is  taken.  The  water  polluted  will,  on  analysis,  prove 
to  contain  compounds  of  nitrogen  which  have  their  origin  in  putres- 
cent  animal  substances  ;  the  quantity  of  these  compounds  diminishes 
as  the  water  is  exposed  to  the  oxidizing  action  of  the  atmosphere ; 
yet,  in  the  water  thus  purified  by  nature,  there  may  be  a  large  number 
of  living  micro-organisms  which  may  cause  disease,  so  that  a  drinking 
water  should  be  examined  not  only  with  the  purpose  of  ascertaining 
the  quantity  of  decaying  animal  substance  present,  but  also  as  to  the 
number  and  kind  of  micro-organisms  contained  in  it.  Water  may 
be  contaminated  by  a  considerable  quantity  of  sewage  and  yet  be 
harmless,  for  it  may  contain  no  harmful  disease  germs ;  while,  on 
the  other  hand,  water  considered  as  pure  may  be  extremely  danger- 
ous, by  reason  of  germs  contained  therein. 

In  concluding  the  chapter  on  this  most  important  subject,  it 
is  well  once  more  to  call  attention  to  the  necessity  of  thoroughly 
understanding  the  changes  of  energy  which  take  place  in  the  forma- 


46  WATER;  REVIEW  or  CHEMISTRY  OF. 

tion  and  decomposition  of  water.  We  must  remember  that  hydro- 
gen and  oxygen  (a  metal  and  a  not-metal)  unite  most  readily, 
provided  some  impulse  (such  as  an  electric  spark  or  fire)  is  added 
to  a  mixture  of  the  gases,  while,  during  the  union,  much  heat  is 
given  off,  so  that  water  possesses  much  less  energy  than  the  ele- 
ments hydrogen  and  oxygen;  and  furthermore  the  application  of 
just  as  much  kinetic  energy  is  necessary  to  decompose  water  as  has 
been  given  off  in  its  formation.  When  the  energy  is  applied  by 
means  of  an  electric  current,  two  volumes  of  hydrogen  and  one  of 
oxygen  are  produced. 

We  have,  by  this  time,  become  somewhat  more  intimately 
acquainted  with  chemical  equations.  The  pupil  should  practise 
writing  the  chemical  equations  which  have  been  given,  until  he  is 
entirely  familiar  with  them,  but  he  must  always  remember  that  the 
mere  memorizing  of  such  equations  is  useless;  that  the  reasons 
why  any  reaction  should  take  place,  and  why  certain  substances  are 
formed  from  any  reaction,  are  of  infinitely  more  importance  than 
the  equations  which,  however,  fix  these  reasons  in  the  memory. 


OZONE;  HISTORY.  41 


CHAPTER   V. 

OZONE   AND  HYDROGEN  DIOXIDE. 

WHEX  compared  with  the  ^extraordinary  chemical  activity  of 
oxygen  at  higher  temperatures,  the  tendency  of  that  element  to 
unite  with  other  substances  under  ordinary  conditions  is  not  very 
marked.  Iron,  copper,  and  similar  substances,  which  are  oxidized 
when  heated  in  the  gas,  remain  unchanged  in  perfectly  pure  oxygen 
at  ordinary  temperatures,  and  we  can  readily  see  that,  did  oxygen 
possess  the  capability  of  oxidizing  under  these  circumstances,  there 
would  result  a  complete  alteration  of  the  existing  conditions  upon 
the  surface  of  the  earth.  The  element  is  not,  however,  limited  to 
the  one  form  which  we  have  discussed ;  it  is  also  capable  of  existing 
in  another  character,  in  which  its  properties  differ  most  remarkably 
from  those  of  ordinary  oxygen. 

An  element  capable  of  existing  in  two  or  more  different  physical 
and  chemical  forms  is  said  to  possess  the  property  of  allotropism ; 
and  the  different  modifications  of  the  same  element  are  called  its 
allotropic  forms.  Oxygen,  then,  exists  in  two  allotropic  forms, 
oxygen  and  ozone ;  in  the  former  one  it  has  the  properties  which 
have  already  been  described ;  in  the  latter  it  possesses  a  marked 
odor;  when  large  quantities  are  present,  it  is  irritating  to  the 
mucous  membrane  of  the  throat  and  nose,  and  it  is  a  most  active 
oxidizer  even  at  ordinary  temperatures.  We  are  acquainted  with 
ozone  only  when  it  is  diluted  with  oxygen  or  air. 

The  fact  that  a  room,  in  which  a  powerful  generator  of  static 
electricity  is  in  action,  becomes  filled  with  an  odor  resembling  that 
of  phosphorus  has  been  known  for  some  time,*  and,  in  1840,  Schon- 
bein  found  that  the  same  odor  is  produced  when  moist  phosphorus 
is  exposed  to  the  atmosphere.  The  substance  causing  this  odor 
was,  at  first,  owing  to  its  resemblance  to  chlorine,  supposed  to  be 
an  element ;  at  a  later  date  investigation  seemed  to  show  that  it 
was  simply  an  oxide  of  hydrogen  differing  from  water,  but,  finally> 

*  Since  1785  (Von  Marum). 


48  OZONE;  RELATION  TO  OXYGEN  BY  VOLUME. 

de  la  Rive  proved  that  if  oxygen,  perfectly  pure  and  dry,  is  passed 
through  a  narrow  glass  tube  in  which  are  inserted  two  platinum 
wires,  between  which  electric  sparks  are  passing,  a  quantity  of 
ozone  is  generated.  This  proved  without  a  doubt  that  ozone  was 
generated  from  oxygen  alone,  and  subsequently  the  proof  was 
brought  that  ozone,  on  heating,  yields  no.thing  but  oxygen. 

Having  discovered  that  ozone  is  simply  oxygen  in  another 
form,  there  remained  to  be  .decided  whether,  in  forming  the  former 
substance  from  the  latter,  any  change  in  the  bulk  of  the  gas 
occurred.  Further  study  showed  that  a  diminution  in  volume 
takes  place;  and  this  contraction  was  such  that  from  3  c.c.  of 
oxygen  there  result  2  c.c.  of  ozone,  and  conversely  from  two  of 
ozone  there  are  formed  three  of  oxygen.  We  have  learned  that  in 
equal  volumes  of  gases  there  are  equal  numbers  of  particles ;  it  fol- 
lows that  if  we  were  able  to  obtain  pure  ozone,  there  would  be  as 
many  particles  of  ozone  in  a  given  volume  as  there  would  be  in  the 
same  bulk  of  oxygen.  Now,  we  have  seen  that  in  the  formation 
of  ozone,  oxygen  contracts  from  3  volumes  to  2,  it  follows  that  a 
given  weight  of  ozone  occupies  only  two-thirds  the  volume  of  the 
same  weight  of  oxygen ;  hence,  the  weights  of  equal  volumes  of 
oxygen  and  ozone  must  be  to  each  other  as  2:3;  and  hence,  if 
there  are  the  same  number  of  molecules  in  equal  volumes  of  each 
gas,  the  weight  of  a  molecule  of  oxygen  must  be  to  that  of  ozone  as 
2  :  3.  We  will  :learn,v  empirically  for  the  present,  that  the  molecule 
of  oxygen  is  composed  of  two  atoms,  and  that  its  molecular  weight, 
as  it  is  the  sum  of  the  atomic  weights  of  the  atoms  composing  the 
molecule,  must  be  32,  it  follows  that  the  molecular  weight  of  ozone 
is  48,  and  that,  if  ordinary  oxygen  has  a  molecule  composed  of  two 
atoms  of  oxygen,  ozone  must  have  one  consisting  of  three.  In  this 
case,  then,  the  cause  of  allotropism  is  evidently  found  in  the  differ- 
ent molecular  structure  of  the  two  modifications  of  the  same 
element,  and  from  this  the  student  will  see  that  a  change  in  the 
composition  of  a  molecule  brings  with  it  a  change  in  the  character 
of  the  substance,  regardless  of  whether  that  molecule  is  composed 
of  atoms  of  the  same  kind  or  of  those  of  different  kinds.  The  two 
reactions, 

S  +  02  =  S02 ,  Sulphur  dioxide, 
0  +  02  =  002  Ozone, 
will  serve  to  make  this  meaning  more  clear.     By  oxidizing  sulphur 


OZONE;  PROPERTIES. 

we  obtain  sulphur  dioxide,  a  body  differing  in  properties  from  both 
sulphur  and  oxygen ;  by  oxidizing  oxygen  we  obtain  ozone,  a  body 
differing  in  properties  from  oxygen,  but  not,  perhaps,  as  markedly 
as  sulphur  dioxide  does  from  sulphur.  An  addition  of  energy  is 
necessary  in  order  to  produce  ozone  from  oxygen ;  it  is  therefore  an 
endothermic  compound,  and,  as  a  consequence,  has  a  great  tendency 
to  break  down  with  the  evolution  of  heat.  The  fact  that  it  can 
oxidize  metals  under  ordinary  conditions,  has  already  been  alluded 
to ;  *  it  also  can  oxidize  a  great  many  organic  substances,  such  as 
albumen,  milk,  shavings,  corks,  or  india  rubber ;  if  such  substances 
are  placed  in  oxygen  containing  ozone,  the  odor  of  the  latter 
disappears  at  once.  As  ozone  is  formed  in  a  great  variety  of  ways, 
for  instance,  by  the  evaporation  of  liquids  or  by  discharges  of  elec- 
tricity, it  follows  that  more  or  less  of  the  substance  must  occur  in 
the  atmosphere  at  times,  but,  owing  to  the  presence  of  oxidizable 
substances,  we  should  scarcely  expect  any  ozone  to  be  present  in 
the  air  of  cities,  t  Large  quantities  of  ozone  would  undoubtedly  be 
harmful  if  inhaled ;  it  has  never  been  proved  that  small  quantities 
have  any  effect.  J 

Ozone  is  a  gas  which  has  a  blue  tint,  which  can  be  seen  by  look- 
ing in  the  direction  of  a  white  paper,  through  a  long  tube  contain- 
ing ozone.  If  it  could  be  obtained  pure,  it  would  undoubtedly  be 
easily  condensed  to  a  liquid,  for,  although  it  is  always  greatly 
diluted  with  oxygen,  it  nevertheless  forms  an  indigo-blue  liquid  at 
temperatures  above  those  required  to  liquefy  oxygen.23 

Hydrogen  and  oxygen  form  two  distinct  compounds,  in  one  of 
which  (water)  there  are  two  parts  by  weight  of  hydrogen  united  to 
sixteen  of  oxygen ;  in  the  other,  two  of  hydrogen  to  thirty-two  of 
oxygen ;  the  existence  of  these  two  compounds  forms  an  excellent 
example  of  the  law  of  multiple  proportions.  We  have  already 
decided  that  the  formula  of  water  is  H20,  and  hence  we  must 
assign  the  formula  H2  O2  to  hydrogen  dioxide,  remembering  that,  as 
we  have  not  been  able  to  obtain  this  substance  in  the  form  of  a  gas, 

*  Some  pure  and  bright  mercury  shaken  with  gas  which  contains  even 
traces  of  ozone,  is  instantly  oxidized,  the  mercury  losing  its  lustre  and,  in 
part,  adhering  to  the  sides  of  the  flask. 

t  The  presence  of  ozone  in  the  atmosphere  has  recently  been  denied.  Ilos- 
vay-Ilosva;  Bulletin  Soc.  Chim.;  [3]  2,  377. 

+  Labbe  and  Oudin;  Comptes  Rendus;  113,  141. 


50  HYDROGEN  DIOXIDE;   PREPARATION. 

H2  02  can  represent  only  the  formula  weight ;  the  molecular  weight 
may  be  any  multiple  of  this  formula  weight,  or  (n  H2  02).* 

(Hydrogen  dioxide  is  prepared  by  adding  a  dilute  acid,  preferably 
sulphuric  acid,  to  barium  dioxide. 

Ba02  +H2S04          =BaS04  +H202, 

Barium  Dioxide  -f-  Sulphuric  acid  =  Barium  sulphate  +  Hydrogen  dioxide. 
Barium  sulphate  is  insoluble  in  water ; 24  it  can  therefore  be 
allowed  to  settle  to  the  bottom  of  the  vessel  in  which  the  dioxide 
of  hydrogen  is  prepared,  and  the  clear  supernatent  liquid  then 
poured  off;  by  allowing  the  excess  of  water  to  evaporate,25  there 
remains  a  very  concentrated  solution  of  hydrogen  dioxide,  having  a 
specific  gravity  of  1.45,  and  which  does  not  freeze  at  —  30°.  The 
concentrated  solution"  must  be  preserved  in  ice,  for  on  warming  to 
ordinary  temperatures,  a  rapid  evolution  of  oxygen  takes  place,  and 
nothing  but  water  remains;  a  too  rapid  heating  of  the  liquid  to 
the  boiling  point  of  water  will  even  cause  it  to  explode.  Dilute 
solutions  of  the  dioxide  have  a  bitter  taste ;  they  are  used  for 
purposes  of  bleaching. 

Hydrogen  dioxide  owes  its  chief  value  to  the  readiness  with 
which  it  yields  its  oxygen,  it  resembling  ozone  in  that  particular ; 
indeed,  one  of  the  most  delicate  tests  for  both  of  these  substances 
is  the  same,  and  owes  its  value  to  their  oxidizing  power.  Hydrogen 
dioxide  or  ozone,  when  either  is  added  to  a  solution  of  potassium 
iodide,  yields  iodine  :  — 

H202  +  2KI  =2KOH  +  21. 

Hydrogen  dioxide  +  Potassium  iodide  =  Potassium  hydroxide  +  Iodine. 
03      +2KI  +H20  =  2KOH  +  02        +21 

Ozone  +  Potassium  iodide  +  Water  =  Potassium  hydroxide  -f  Oxygen  +  Iodine. 

In  the  latter  case  the  addition  of  water  is  necessary;  in  the 
former  a  compound  of  hydrogen  and  oxygen  which  yields  water  is 
already  present,  but  in  both  cases  the  oxygen  given  off  changes  the 
iodide  to  the  hydroxide  of  potassium.  Iodine  has  the  power  of 
turning  starch  paste  to  a  deep  blue  color,  so  that  the  addition  of 
some  of  this  substance  will  render  even  minute  traces  of  iodine 
visible. )  > 

*  The  molecular  weight  of  hydrogen  dioxide  is  probably  34,  and  the  mole- 
cule H2O2,  a  fact  which  has  been  discovered  by  determining  the  freezing  point 
of  hydrogen  dioxide  solutions.  See  Orndorff  and  White;  Zeitschrift  fiirPliysik. 
Chemie;  12,  1. 


NASCENT   STATE.  51 

Both  ozone  and  hydrogen  dioxide  owe  their  peculiar  powers  of 
oxidation  to  the  fact  that  they  can  yield  oxygen  in  a  condition 
known  as  the  nascent  state.  Oxygen,  as  well  as  a  number  of  other 
elements,  exists  as  molecules,  each  molecule  being  formed  of  two 
atoms.  A  considerable  amount  of  energy  is  necessary  to  decompose 
these  molecules ;  indeed,  in  the  case  of  hydrogen,  for  instance,  it  is 
doubtful  whether  any  heat  which  we  can  command  will  be  able  to 
decompose  its  molecules  into  individual  atoms.  As  a  consequence, 
it  follows  that  these  atoms  possess  much  more  chemical  energy  than 
do  the  molecules,  and  hence  must  have  a  greater  tendency  to  unite 
with  some  other  atom  or  with  some  molecule.  When  an  element  is 
liberated  from  any  of  its  compounds,  it  must  at  first  exist  as  indi- 
vidual atoms,  no  matter  how  short  a  time  is  necessary  for  the  atoms 
to  unite  to  form  molecules.  If,  however,  any  substance  is  present 
011  which  the  atoms  can  act,  they  will  primarily  react  with  that  sub- 
stance. It  follows,  therefore,  that  elements  are  chemically  most 
active  at  the  very  moment  of  their  liberation  from  compounds  (in 
statu  nascendi).  If  we  pass  hydrogen  gas  through  nitric  acid  no 
change  will  take  place,  no  matter  how  long  we  may  continue  the 
operation ;  but  if  we  generate  hydrogen  within  the  acid,  as,  for  in- 
stance, by  placing  a  piece  of  zinc  in  nitric  acid,  the  hydrogen  will 
rob  the  nitric  acid  of  its  oxygen,  forming  water  and  an  oxide  of 
nitrogen  containing  less  oxygen  than  does  nitric  acid.1*  Instances 
of  the  action  of  elements  in  the  nascent  state  are  extremely  numer- 
ous, but  we  are  also  aware  of  a  number  of  cases  where  compounds 
are  more  energetic,  chemically,  at  the  moment  of  their  formation 
than  at  any  other  time,  and  in  such  cases  this  explanation  of  the 
nascent  state  is  inadequate.  The  compound  CO,  carbon  monoxide, 
can  act,  under  certain  conditions,  as  if  it  were  in  the  nascent  state ; 
but  we  have  no  reason  to  suppose  that  this  compound  ever  exists 
otherwise  than  as  the  molecule  represented  by  the  formula  CO. 
The  above  explanation  of  the  nascent  state,  if  correct,  is  probably 
applicable  only  in  a  limited  number  of  cases.  The  fact  remains, 
however,  no  matter  what  reason  we  see  fit  to  assign  to  the 
phenomenon,  that  elements  frequently  enter  into  reaction  at  the 
moment  of  their  liberation  from  compounds,  where  they  would  be 
entirely  indifferent  under  other  circumstances.  Oxygen  in  the 
nascent  state  is  liberated  by  hydrogen  dioxide  and  by  ozone,  for 

*  See  chapter  xxvi. 


52  NASCENT    STATE. 

0 

one  molecule  of  ozone,  |  >  0  (03)  breaks  down  into  one  molecule  of 
0 

0  0 

oxygen  and  an  atom  of  the  same  element    |  >  0  =    |    +0,  while 

0  0 

the  same  quantity  of  hydrogen  dioxide  decomposes  into  a  molecule 

H-0      H        +0 
of  water  and  an  atom  of  oxygen,        |  =    >  0         ,  and,  as  a  coii- 

H-0      H 

sequence,  both  ozone  and  hydrogen  dioxide  are  powerful  oxidizers 
and  bleachers.  The  tendency  of  oxygen  to  unite  with  other  atoms, 
when  it  is  liberated  from  ozone,  is  so  great  that  it  can  even  take 
oxygen  away  from  other  compounds  to  form  a  molecule  of  oxygen  ; 
for  instance,  -when  ozone  is  brought  in  contact  with  silver  oxide, 
the  following  reaction  takes  place  :  — 


and  similar  changes  are  produced  with  hydrogen  dioxide.* 

*  The  occurrence  of  hydrogen  dioxide  in  the  atmosphere  is  doubtful. 
See  Schone  ;  Berichte  d.  Deutsch.  Chem.  Gesell.;  26,  3011.  Hydrogen  dioxide 
has  lately  been  obtained  in  a  pure  state  by  distillation  in  a  vacuum  ;  its  boiling 
point  is  85°  under  68  mm.  pressure.  R.  Wolfenstein,  Ber.  d.  Deutsch.  Chem. 
Gesell.;  27,  3307. 


THE   HALOGENS  ;    COMPARISON    OF.  53 


CHAPTER   VI. 

THE    HALOGENS. 

WE  have  now  studied  a  metal,  a  typical  not-metal,  and  the  pro- 
duct formed  by  the  union  of  the  two ;  and  have  gained  an  insight 
into  quite  a  number  of  chemical  reactions,  as  well  as  into  the  mode 
of  action  of  the  molecules  and  atoms ;  so  that  we  will  now  go  to  the 
discussion  of  the  elements  by  families,  taking  them  up  in  the 
natural  order  assigned  to  them  by  their  atomic  weights,  remember- 
ing that,  as  was  said  in  the  introduction,  the  properties  of  the 
various  elements  are  determined  by  their  atomic  weights.  The  first 
group  of  elements  which  we  will  study  includes  those  which  are 
most  not-metallic  in  their  characteristics,  and  the  plan  will  be  to 
work  from  this  family  through  others  with  a  diminishing  not- 
metallic  character,  until  finally  we  arrive  at  groups  composed 
entirely  of  metals. 

The  group  of  Halogens  (salt  producers)  comprises  four  elements : 

Fluorine,  atomic  weight  19. 

Chlorine,       «  «       35.45.  ' 

Bromine,       "  "       79.95.  <%  0 

Iodine,  "  «     126.85.  ,  ., 

With  increasing  atomic  weight  we  have  a  decrease  in  the  not-metallic 
properties  of  the  elements  included  in  this  family,  and  this  change 
is  well  shown  in  the  decreasing  stability  of  the  compounds  of  the 
halogens  with  the  metals.  If  we  examine  the  hydrogen  compounds, 
which  are  formed  by  the  union  of  one  atom. of  hydrogen  with  one  of 
the  halogen  (HF,  HC1,  HBr,  HI),  we  are  at  once  impressed  by  this 
change  —  for  hydroiodic  acid  (HI)  decomposes  most  readily  upon 
heating,  a  hot  wire  introduced  into  the  gas  will  change  it  to  hydro- 
gen and  iodine ;  hydrobromic  acid  (H  Br)  is  less  readily  separated 
into  its  elements ;  hydrochloric  acid  (H  Cl)  is  broken  down,  if  at 
all,  only  by  the  application  of  a  very  great  amount  of  heat ;  and  we 
have,  so  far  as  we  know,  never  effected  a  decomposition  of  hydro- 


54  THE   HALOGENS  ;   COMPARISON   OF. 

fluoric  acid  by  heat  alone.  We  can  see  this  same  difference  illus- 
trated by  bringing  chlorine  into  a  solution  containing  sodium 
"bromide  or  iodide;  the  chlorine  will  at  once  form  sodium  chloride, 
liberating  bromine  or  iodine ;  while  bromine,  added  to  a  solution  of 
sodium  iodide,  will  set  iodine  free,  and  fluorine  would,  without 
doubt,  decompose  the  compounds  of  any  of  the  other  halogens  with 
the  metals.  This  difference  in  the  character  of  the  halogens  is 
shown  by  the  heat  of  formation  of  the  hydrogen  compounds  given 
in  the  table  at  the  end  of  Chapter  XI.  With  increasing  atomic 
weight  we  necessarily  have  changes  in  the  physical  properties  of 
the  elements.  Fluorine  is  a  nearly  colorless  gas,  chlorine  a  yellow- 
ish-green gas,  rather  easily  converted  to  a  liquid,  bromine  is  a  dark 
brown  liquid  at  ordinary  temperatures,  while  iodine  is  a  solid  of 
almost  metallic  appearance.  The  halogens  all  have  a  peculiar  odor, 
and  attack  the  skin  and  mucous  membrane.  Other  points  of  resem- 
blance will  become  apparent  in  the  detailed  study  of  the  elements. 
They  all,  with  the .  exception  of  fluorine  and  bromine,  form  oxides, 
and,  all  but  fluorine,  acids  with  oxygen  and  hydrogen ;  a  study  of 
the  latter  will  be  deferred  until  a  subsequent  chapter  is  reached. 
The  halogens  themselves  are  never  found  as  free  elements,  but  occur 
united  to  some  metal,  as  the  fluoride,  chloride,  bromide,  or  iodide. 
The  metals  most  frequently  found  united  with  the  halogens  are 
sodium,  potassium,  magnesium,  or  calcium,  so  that  for  instance, 
sodium  chloride  (common  salt)  is  the  most  frequently  occurring 
compound  of  chlorine.  Having  given  a  few  of  these  general  char- 
acteristics, we  will  go  to  the  study  of  the  individual  elements. 


FLUORINE;  OCCURRENCE,  PREPARATION.  55 


CHAPTER   VII. 

FLUORINE    AND    HYDROFLUORIC    ACID. 

Symbol,  F;  atomic  weight,  19.0^ Formula,  HF;  specific  gravity  of 
liquid,  .9879 ;  of  gas,  air  =  1,  is  1.364 ;  H2  =  2,  is  39.32 ;  the 
molecule  is  represented  by  the  formula  H2  F2. 

THIS  element  chiefly  occurs  in  nature,  combined  with  the  metal 
•calcium,  as  fluorspar  (fluorite,  CaF2),  a  crystalline  mineral  quite 
frequent  of  occurrence ;  in  addition  to  this,  cryolite,  a  fluoride  of 
sodium  and  aluminium  (A1F8 ,  3  NaF),  occurs  in  large  masses  in 
Greenland,  and  is  a  considerable  source  of  fluorine  compounds  ; 
while  small  quantities  of  fluorides  occur  in  the  enamel  of  teeth  and 
blood,  and  traces  of  the  same  are  found  in  sea-water. 

Fluorine  has,  because  of  its  great  chemical  affinity  for  other  ele- 
ments, be  they  metal  or  not-metal,  until  recently  resisted  all  attempts 
to  isolate  it ;  this  is  due  to  the  fact  that  it  would  combine  with  other 
substances  as  soon  as  liberated  from  its  compounds.  Quite  recently 
a  French  chemist,  Moissan,  succeeded  in  preparing  the  element  by 
electrolysis  of  perfectly  pure  liquid  hydrofluoric  acid  placed  in  a  U 
shaped  platinum  tube  and  cooled  to  a  low  temperature.*  When  the 
electric  current  passes  through  hydrofluoric  acid  the  latter  is  decom- 
posed into  hydrogen  and  fluorine,  just  as  we  decompose  water  into 
hydrogen  and  oxygen,  the  hydrogen  separating  at  the  negative  pole, 
the  fluorine  at  the  positive. 

The  element  is  a  pale  yellow  gas,f  which  does  not  attack  plati- 
num at  low  temperatures,  but  instantly  unites  with  elements  such 
as  silicon,  boron,  arsenic,  sulphur,  iodine,  iron ;  and  with  organic  sub- 
stances, such  as  cork,  petroleum,  etc.  The  substances  so  attacked 
take  fire  in  the  gas,  so  that  all  the  phenomena  of  combustion  in 
oxygen  are  repeated  with  fluorine  in  a  more  violent  degree  and 
under  ordinary  conditions.  If  the  gas  is  passed  into  water,  the 

*  By  liquid  methyl-chloride,  boiling  point  —  22.5°. 
t  Moissan;  Comptes  Rendus;  109,  937. 


56         FLUORINE;  PROPERTIES.    HYDROFLUORIC  ACID. 

latter  is  instantly  decomposed,  ozone  *  and  hydrofluoric  acid  being 
produced.  This  reaction  is  very  interesting,  for  the  power  of 
decomposing  water  which  members  of  the  halogen  family  possess, 
diminishes  with  increasing  atomic  weight.  The  reaction  is  as 

follows  :  — 

2  F       +  Ho  0  =  2  HF  +  O. 

Fluorine  +  Water  =  Hydrofluoric  acid  +  Oxygen. 

The  atoms  of  nascent  oxygen  when  liberated  can  combine  with 
each  other  to  form  ozone. 

The  compound  of  hydrogen  and  fluorine,  hydrofluoric  acid,  was 
first  identified  as  a  peculiar  acid  by  Scheele  (1771),  although  the 
fact  that  a  substance  which  would  attack  glass  could  be  produced 
by  the  action  of  sulphuric  acid  on  fluorspar  had  been  known  for 
some  time  (since  1670).  The  nature  of  hydrofluoric  acid  at  first 
was  misunderstood,  owing  to  the  opinion  formerly  held  by  chemists 
that  all  acids  must  contain  oxygen,  so  that  the  existence  of  oxygen 
was  presupposed  in  hydrofluoric  acid.  We  now  know  that  no  com- 
pounds of  fluorine  and  oxygen  exist. 

Hydrofluoric  acid  can,  as  we  have  seen,  be  produced  by  the 
action  of  fluorine  on  water  and  by  direct  union  of  hydrogen  and 
fluorine,  just  as  water  is  formed  by  direct  union  of  hydrogen  and 
oxygen.  So  great  is  the  tendency  to  form  hydrofluoric  acid  that 
fluorine  will  probably  take  hydrogen  away  from  any  other  com- 
pound containing  that  element.  To  prepare  hydrofluoric  acid  for 
the  laboratory  or  for  commercial  purposes,  other,  less  expensive, 
means  than  the  ones  which  have  been  given  are  resorted  to.  If  a 
fluoride,  such  as  sodium  fluoride  or  calcium  fluoride,  is  treated  with 
sulphuric  acid,  the  following  reaction  takes  place  :  f  — 

2NaF  +  H2S04  =  Na2S04  -f  2  HF 

Sodium  fluoride    +  Sulphuric  acid  =  Sodium  Sulphate  +  Hydrofluoric  acid. 

CaF2  +H2S04  =  CaS04  +  2  HF 

Calcium  fluoride  +  Sulphuric  acid  =  Calcium  Sulphate  +  Hydrofluoric  acid. 

Reactions  such  as  the  above  are  very  frequently  met  with  ;  the 
hydrogen  of  the  acid  simply  exchanging  places  with  the  metal  of 

*  If  a  tube,  filled  with  fluorine,  has  a  few  drops  of  water  admitted  to  it, 
ozone  is  formed  in  such  quantity  that  the  contents  are  temporarily  changed  to 
a  deep  blue  color. 

t  As  hydrofluoric  acid  attacks  glass,  the  pure  substance  must  be  prepared 
in  platinum  vessels. 


HYDROFLUORIC   ACID;   PROPERTIES.  57 

the  salt  to  form  a  new  salt  and  a  new  acid.  Such  reactions  are 
designated  as  double  decompositions,  and  as  a  practical  hint  we  can 
say  that  they  take  place  when  an  insoluble  or  a  volatile  substance 
can  be  produced.*  We  shall  inquire  more  closely  into  the  nature 
of  double  decompositions  when  we  have  studied  a  larger  number  of 
chemical  reactions.  The  method  given  above  is  one  which  is  very 
frequently  employed  in  the  preparation  of  the  various  acids  from 
their  salts. 

Hydrofluoric  acid  is  a  colorless,  mobile  liquid  which  freezes  at 
—  102°,  boils  at  -f-  19°  (about  the  temperature  of  a  warm  room), 
and  which  fumes  strongly  in  the  air  because  of  its  attraction  for 
moisture.  The  vapor  of  hydrofluoric  acid  is  very  irritating  even 
when  inhaled  in  small  quantities,  while  any  considerable  amount 
can  cause  death.  A  drop  of  the  acid  put  on  the  hand  causes  a  most 
painful  blister,  which  ultimately  changes  to  a  slowly  healing  ulcer, 
so  that  great  care  must  be  exercised  in  handling  this  acid.  The 
usual  commercial  acid  is  a  solution  of  hydrofluoric  acid  in  water. 
It  is  a  colorless,  extremely  acid  liquid,  it  fumes  in  the  air,  and  is 
transported  in  bottles  made  either  of  paraffin  or  of  guttapercha, 
for  the  acid  readily  attacks  glass. 

*  Hydrofluoric  acid  is  a  volatile  substance. 


58  CHLORINE;  OCCURRENCE,  HISTORY. 


CHAPTER   VIII. 

CHLORINE. 

5a 

Symbol,  Cl ;  atomic  weight,  35.45  ;  specific  gravity,  air  =  1  is,  2.46, 

H2  =  2,  is  70.84.     1  c.c.  iveighs  .0031825  gram. 

CHLORINE  occurs  in  nature  combined  with  various  metals  as 
chlorides,  never  as  the  free  element.  The  most  important  chloride 
is  that  of  sodium  (Nad,  sodium  chloride,  or  common  salt).,;  This 
substance  forms  the  major  portion  of  the  solid  residue  left  upon 
evaporation  of  sea-water,  and  is  consequently  the  larger  part  of  the 
salt  beds  of  marine  origin  and  of  those  composed  of  rock  salt. 
Chlorine  also  occurs  in  considerable  quantities  combined  with 
potassium  as  the  mineral  sylvin  (KC1,  potassium  chloride),  and  as 
a  chloride  of  magnesium  and  of  potassium  ( K  Cl,  Mg  C12  +  6  H2  0, 
called  carnallite)  ;  the  chlorides  of  iron,  lead,  silver,  etc.,  occur  in 
small  quantities,  while  chlorides  are  found  in  the  tissues  and  fluids 
of  plants  and  animals  and  in  their  ashes. 

The  element  was  not  discovered  until  1774,  at  the  beginning  of 
the  period  in  which  the  great  advances  in  chemistry  recorded  in 
the  introductory  chapter  were  made ;  the  first  chemist  to  prepare 
chlorine  was  Scheele.  He  called  it  dephlogisticated  muriatic  acid, 
for  it  was  muriatic  acid  from  which  phlogiston  (hydrogen)  had  been 
extracted.  Chlorine  was  for  some  time  supposed  to  be  a  compound 
of  oxygen  with  an  unknown  element  called  murium ;  Sir  Humphry 
Davy  first  definitely  asserted  that  chlorine  was  an  element,  naming 
the  element  chlorine  from  x\wp6s,  greenish  yellow,  the  gas  having 
that  color. 

1.  Preparation  of  Chlorine  by  Electrolysis. 

Evidently,  in  order  to  prepare  chlorine,  our  method  must  be  to 
remove  the  metallic  constituent  from  some  chloride ;  hydrogen  chlo- 
ride (hydrochloric  acid)  being  the  chloride  easiest  available.  If  we 
subject  a  concentrated  solution  of  hydrochloric  acid  to  the  action  of 
an  electric  current,26  we  shall  decompose  the  substance  in  exactly 
the  same  manner  as  we  can  hydrofluoric  acid,  with  the  difference 


CHLORINE;  PREPARATION.  59 

that  we  can  readily  perform  this  operation  in  glass  vessels  (which 
are  not  attacked  by  chlorine),  as  the  chemical  energy  displayed  by 
chlorine  is  not  so  great  as  that  of  fluorine.  We  can  also  remove 
hydrogen  from  hydrochloric  acid  by  other  means  ;  for  instance, 
by  some  oxidizing  agent,  when  water  and  chlorine  are  produced,  as 
follows  :  — 

2HC1  +0         =H20+2C1.   • 

Hydrochloric  acid  +  Oxygen  =  Water  +  Chlorine. 

The  oxygen  of  the  atmosphere  can  accomplish  this  under  proper 
conditions,  and  a  process  of  commercial  preparation  of  chlorine* 
has  its  origin  in  this  fact.  If  a  mixture  of  hydrochloric  acid  and 
oxygen  is  passed  through  a  heated  tube  in  which  are  placed  pieces 
of  porcelain  or  fire-brick,  saturated  with  a  solution  of  copper  sul- 
phate, chlorine  and  water  are  formed.  The  copper  sulphate  re- 
mains unchanged  at  the  end  of  the  reaction,  so  that  the  reason  for 
its  peculiar  action  is  not  understood. 

2.  Preparation  of  Chlorine  from  Manganese  Dioxide  and  Hydro- 
chloric Acid. 

Manganese  dioxide  is  the  most  convenient  oxidizer  for  the 
preparation  of  chlorine ;  when  it  is  brought  in  contact  with  hydro- 
chloric acid,  the  following  reaction  takes  place  : 27  — 

1.  Mn02  +4HC1  =MnCl2  +2HaO+2Cl. 

Manganese  dioxide .+  Hydrochloric  acid=Manganous  chloride  +  Water  +  Chlorine. 

The  manganese  dioxide  furnishes  the  oxygen  which  changes 
the  hydrogen  of  hydrochloric  acid  to  water,  and  at  the  same  time 
in  all  probability  a  chloride  of  manganese,  having  four  atoms  of 
chlorine  in  the  formula  weight,  is  formed,  thus :  — 

2. 

This  chloride  is,  however,  very  unstable  and  breaks  down  into 
manganous  chloride  and  chlorine,  as  follows :  — 

so  that,  combining  equations  2  and  3,  we  obtain  equation  1.  As  a 
rule,  chemical  equations  are  expected  only  to  represent  the  ultimate 

*  Deacon's  process. 


60  CHLORINE;   PREPARATION. 

product  of  any  chemical  reaction,  as  does  equation  1,  but  if  we 
wish  thoroughly  to  understand  chemical  changes  we  must  not  be 
contented  with  the  mere  equation,  but  should  also  inquire  into  all 
of  the  intermediary  stages  which  bring  about  chemical  reactions. 
Other  oxidizing  agents,  as  well  as  manganese  dioxide,  *  are  capable 
of  furnishing  oxygen  to  form  water  and  chlorine  from  hydro- 
chloric acid. 

3.  Preparation  of  Chlorine  from  Sodium  Chloride,  Sulphuric  Acid, 
and  Manganese  Dioxide. 

A  mixture  of  common  salt  (Na  Cl,  sodium  chloride)  and  sulphuric 
acid  yields  hydrochloric  acid,  so  that  it  is  often  convenient  to  pre- 
pare chlorine  by  mixing  sodium  chloride  and  manganese  dioxide  in 
a  flask  and  then  adding  sulphuric  acid.  In  this  case  it  must  be 
remembered  that  the  manganous  chloride  (MnCl2)  which  might 
be  formed  (equation  3,  preceding  page,)  would  also  be  acted  upon 
by  the  sulphuric  acid,  forming  manganous  sulphate  and  hydro- 
chloric acid,  which  would  further  be  converted  by  the  manganese 
dioxide,  according  to  equation  2  :  — 

Mn  C12  +  H2  S04      =  Mn  S04  +  2  H  Cl. 

MnO2  +  2  H2S04  +  MnCl2  =  2  MnS04+2  H20+2  Cl; 

so  that,  when  the  usual  mixture  of  sodium  chloride  and  manganese 
dioxide  is  used,  the  following  complete  reaction  takes  place :  — 

Mn02  +  2  Na  Cl  +  2  H2S04  =  MnS04  +  Na2S04  +  2  H20  +  2  Cl. 

In  all  cases  the  principle  of  the  reaction  is  the  one  given  on  the 
previous  page,  the  sodium  chloride  serving  simply  to  furnish  hydro- 
chloric acid.  ) 

4.   Other  Methods  for  Preparing  Chlorine. 

Several  other  methods  for  the  preparation  of  chlorine  have  been 
devised ;  some  of  these  are  of  commercial  value,  as,  for  instance, 
the  preparation  by  heating  magnesium  chloride  in  a  current  of  air, 

*  Such  oxidizing  agents  are  potassium  permanganate  (K  Mn  O4),  potas- 
sium bichromate  (K2Cr2O7),  nitric  acid  (HNO3),  etc.  In  all  these  cases  the 
principle  is  the  same,  the  sole  object  being  to  remove  hydrogen  from  hydro- 
chloric acid,  forming  chlorine  and  water.  Manganese  dioxide  is  employed 
because  it  is  cheap. 


CHLORINE;    PROPERTIES.  61 

the  reaction  resembling  that  of  the  action  of  oxygen  on  hydrochloric 

acid :  — 

MgCl2  +0          =  MgO  +  2C1 

Magnesium  chloride  +  Oxygen  =  Magnesium  oxide  +  Chlorine. 

The  magnesium  oxide  formed  can  be  dissolved  in  hydrochloric  acid, 
once  more  forming  magnesium  chloride,  and  thus  the  process  can 
be  continued  without  serious  loss  of  magnesium.* 

(Chlorine,  at  ordinary  temperatures,  is  a  greenish  yellow  gas,  the 
color  of  which  becomes  darker  upon  heating.  It  has  a  peculiar, 
irritating  odor,  which  must  be  tested  only  when  the  gas  is  very 
dilute,  for  any  great  quantity  of  chlorine  entirely  destroys  the  sense 
of  smell,  causing  inflammation  of  the  mucous  membrane  of  the 
throat  and  lungs,  coughing  and  hemorrhages.  An  annoying 
catarrh  follows  the  inhalation  of  the  gas,  so  that  great  care  must 
be  exercised  in  working  with  chlorine.  The  presence  of  any  con- 
siderable quantity  of  chlorine  in  the  air  may  cause  death.  A  pres- 
sure of  six  atmospheres,  at  0°,  condenses  chlorine  to  a  liquid ;  at 
ordinary  atmospheric  pressure  it  becomes  liquid  at  —  35°  and  it 
freezes  at  — 102° ;  its  specific  gravity,  air  being  one,  is  2.46  at 
temperatures  up  to  1200°,f  at  higher  temperatures  the  specific 
gravity  becomes  less,  being  but  2.02  at  1400° ;  this  shows  that  the 
chlorine  molecules  begin  to  decompose  into  the  individual  atoms  at 
about  white  heat.J  This  decomposition  of  molecules  into  simpler 
ones  or  into  atoms  is  termed  dissociation ;  the  temperature  of  dis- 
sociation varies  with  different  substances  according  to  the  amount 
of  heat  given  off  in  their  formation ;  it  being,  of  course,  necessary 
to  add  as  much  energy  to  decompose  a  substance  as  is  given  off  in 
its  formation.  At  very  high  temperatures,  such  as  exist  in, the 
chromosphere  of  the  sun,  dissociation  of  all  complex  substances 
is  complete,  so  that  in  such  a  situation  no  chemical  compound  is 
possible. 

Chlorine  is  tolerably  soluble  in  water,  one  volume  of  that  liquid 

*  A  convenient  laboratory  method  for  preparing  chlorine  is  by  the  decom- 
position of  pressed  cubes  of  calcium  hypochlorite  (chloride  of  lime)  by  means 
of  diluted  hydrochloric  acid.  (See  chapter  XVIII.) 

t  This,  H  =  2  would  give  a  specific  gravity  of  70.84,  the  atomic  weight  of 
chlorine  is  35.45,  which  would  give  a  molecular  weight  of  C12  =  70.90,  hence 
below  1200°  chlorine  molecules  consist  of  two  atoms  to  the  molecule. 

$  Langer  and  V.  Meyer;  Berichte  d.  Deutsch.  Cheni.  Gesell. ;  15,  1721. 


62 


CHLORINE;  CHEMICAL  REACTIONS. 


absorbing  2.5  volumes  of  chlorine  at  ordinary  temperatures.  The 
solution  of  the  gas  so  produced,  known  as  chlorine  water,  has  the 
odor  of  chlorine  and  many  of  the  chemical  properties  of  the  gas. 
When  chlorine  water  is  cooled  nearly  to  the  temperature  of  freezing 
water  it  is  changed  to  a  transparent,  crystalline  substance  (chlorine 
hydrate,*  having  the  composition  2  Cl  -}-  8  H2  0),  which  slowly 
gives  off  chlorine  at  low  temperatures,  and 
rapidly  upon  heating.  If  a  few  chlorine 
hydrate  crystals  are  placed  in  one  end  of  a 
bent  glass 'tube,  which  is  sealed  at  the  other 
extremity,  cooled  by  snow  and  salt  (such  a 
tube  is  shown  by  Fig.  6),  chlorine  will  con- 
dense to  a  liquid  in  the  cold  part  of  the 
tube  as  soon  as  the  crystals  are  gently 
warmed. 

Chemically,    chlorine    greatly    resembles 
oxygen,  with  this  distinction ;  while  oxygen, 

under  common  circumstances,  Is  inactive,  chlorine,  at  ordinary 
temperatures,  unites  with  many  elements,  metallic  or  not-metallic, 
to  form  chlorides,  the  formulae  of  which  bear  a  great  resemblance 
to  the  corresponding  oxides  ;  for  example  :  — 

P205  PC15  A1208  Aids 

Phosphorus  pentoxide.  Phosphorus  pentachloride.  Aluminium  oxide.  Aluminium  chloride. 


Fig.  8. 


P203 

Phosphorus  trioxide. 


PC13 

Phosphorus  trichloride. 


The  difference  between  the  formulae  of  chlorides  and  oxides,  as 
seen  from  the  above,  is  that  in  the  oxides  one-half  as  many  atoms 
of  oxygen  unite  with  one  atom  of  the  other  element  entering  into 
the  compound  as  do  chlorine  atoms  in  the  chlorides ;  this  relation- 
ship becomes  clearer  if  we  double  the  formulae  of  the  chlorides  for 
purposes  of  comparison  :  — 


P205, 


P2C110. 
P2C16. 


One  oxygen  atom  is  therefore  capable  of  taking  the  place  of 
two  chlorine  atoms  in  chemical  compounds,  and  in  writing  chemical 


Prepared  by  passing  chlorine  into  ice-water. 


CHLOKINE;  COMBUSTION  IN.  63 

formulae  this  difference  must  always  be  borne  in  mind.     What  is 
true  of  chlorine  applies  to  the  other  halogens  as  well. 

The  union  of  the  various  elements  with  chlorine  and  the  resem- 
blance between  these  reactions  and  combustions  in  oxygen  is  made 
clear  by  the  following  examples  :  — 

'Phosphorus,  which  has  previously  been  ignited  in  the  air,  will 
continue  to  burn  in  an  atmosphere  of  chlorine  with  a  pale  greenish 
flame :  — 

P  +  3  Cl  =  P  C18  (Phosphorus    trichloride). 

Pronounced  metals,  such  as  sodium,  when  heated  to  their 
kindling  temperature,  will  burn  in  chlorine  gas :  — 

Na  +  Cl  =  NaCl  (Sodium  chloride). 

Carbon,  when  heated  in  the  presence  of  chlorine,  will  form 
a  chloride :  — 

C  +  4  Cl  =  C  C14  (Carbon  tetrachloride).* 

Chlorine  unites  with  hydrogen  with  such  facility  that  it  will 
even  extract  hydrogen  from  its  numerous  compounds  with  carbon. 
A  piece  of  filter  paper,  saturated  with  turpentine  and  placed  in 
chlorine  gas,  will  take  fire  ((sometimes  with  explosive  violence)}  form- 
ing hydrochloric  acid,  (and  separating,  in  the  form  of  soot,  the 
carbon  which  was  in  the  turpentine^f  But  it  is  not  only  from  com- 
pounds of  carbon  that  chlorine  will  extract  hydrogen ;  it  will  do  so 
from  a  multitude  of  other  substances  containing  this  element ;  for 
example,  chlorine  will  decompose  sulphuretted  hydrogen  ( H2  S), 
ammonia  (NH8),  or  even  water,  in  each  case  forming  hydro- 
chloric acid,  and  liberating  the  element  previously  combined  with 
hydrogen :  — 

H2S  +  2C1      =2HC1  +S, 

Hydrogen  sulphide  +  Chlorine  =  Hydrochloric  acid  +  Sulphur, 
NH3  +3C1      =3HC1  +N, 

Ammonia  +  Chlorine  =  Hydrochloric  acid  +  Nitrogen, 

H20  +  2C1      =2HC1  +0. 

Water  +  Chlorine  =  Hydrochloric  acid  +  Oxygen. 

*  The  pupil  will  note  that  in  this  formula  two  chlorine  atoms  take  the 
place  of  one  oxygen  atom,  as  will  be  seen  by  comparing  the  formulae  C  C14 
and  CO., . 

t  Turpentine  is  a  compound  of  carbon  and  hydrogen. 


64  CHLORINE   WATER. 

The  cause  for  these  reactions  in  each  case  is  that,  in  contact 
with  hydrogen,  chlorine  possesses  greater  chemical  energy  than  the 
three  other  elements.  This  is  shown  by  the  fact  that  hydrochloric 
acid,  when  formed  from  its  elements  and  then  dissolved  in  water, 
has  a  greater  heat  of  formation  than  water,  sulphuretted  hydrogen, 
or  ammonia. 

Thus,  NH3  formed  from  its  elements,  gives  204  K,* 
Ho  S        "          "      «          «  «       73  K, 

H2  0        "          "      «          «  "     684  K, 

2HC1        "          «      «          «  «     786  K.f 

Chlorine  water  placed  in  the  sunlight  will  form  hydrochloric 
acid  and  liberate  oxygen ;  J  but  chlorine  and  water  can  yield  oxygen 
and  form  hydrochloric  acid  even  without  the  aid  of  sunlight,  pro- 
vided some  substance  is  present  which  can  be  oxidized.  It  is  to 
this  property  that  chlorine  owes  its  commercial  value,  its  chief 
industrial  use  being  as  a  bleaching  agent,  its  power  of  bleaching 
depending,  at  least  in  the  great  majority  of  cases,  upon  its  capability 
of  liberating  nascent  oxygen  from  water.  That  this  is  the  case  can 
be  proven  by  placing  a  piece  of  colored  calico  in  a  jar  of  dry  chlo- 
rine j  no  bleaching  will  take  place  until  water  is  added,  and  then 

*  K  stands  for  the  quantity  of  heat  which  a  gram  of  water  loses  when 
cooled  from  100°  to  0°.  In  speaking  of  heats  of  formation  or  of  the  thermal 
changes  during  reactions,  the  quantities  of  substances  reacting  are  taken  at 
as  many  grams  as  are  expressed  by  the  atomic  or  formula  weights.  Thus, 
when  we  say  the  heat  of  formation  of  H2  O  is  680  K,  we  mean  that  two  grams 
of  hydrogen  uniting  with  sixteen  of  oxygen  give  680  K.  By  heat  of  solution 
we  mean  the  heat  given  off  by  dissolving  the  formula  weight  of  a  substance, 
in  grams,  in  an  unlimited  quantity  of  water.  Thus  "  the  heat  of  solution  of 
hydrochloric  acid  is  172  K,"  means  that  36.45  grams  of  hydrochloric  acid  give 
off  172  K  while  dissolving  in. an  unlimited  amount  of  water.  The  unit,  K, 
and  the  figures  which  are  given  here  and  in  subsequent  parts  of  the  work, 
correspond  to  those  employed  by  Ostwald.  See  also  Ostwald's  Outlines  of 
General  Chemistry,  Walker's  translation,  Macmillan,  1890,  p.  212. 

t  These  figures  refer  to  the  heats  of  formation  of  the  compounds  from 
their  elements,  and  in  the  presence  of  an  excess  of  water. 

(t  Chlorine  water  liberates  oxygen  and  forms  pure  hydrochloric  acid  only 
when  the  solution  is  placed  in  the  direct  sunlight.  In  diffused  light,  even  in 
bright  daylight,  in  addition  to  oxygen,  hypochlorous  acid  and  chloric  acid  as 
well  as  hydrochloric  acid  are  produced.  Pedler;  Journal  of  the  Chem.  Soci- 
ety ;  1890,  613/j 


CHLORINE;    OXIDIZING   POWERS.  65 

the  bleaching  action  of  chlorine  at  once  becomes  apparent.28  Chlo- 
rine is  very  frequently  employed  as  an  oxidizing  agent  in  laboratory 
work  ;\  an  oxidizing  agent  being  a  body  which  can  chemically  furnish 
oxygen,  either  per  se,  or  by  the  decomposition  of  some  oxide. )  Ex- 
amples of  the  former  class  are  such  bodies  as  manganese  dioxide, 
nitric  acid,  potassium  bichromate  or  potassium  permanganate  ;  all  of 
these  substances  are  direct  oxidizers,  and,  while  they  oxidize,  they 
themselves  are  reduced.  Examples  of  indirect  oxidizers  are  chlorine, 
bromine,  or  iodine,  for  these  elements  decompose  water  in  order  to 
accomplish  the  same  result. 


66  HYDROCHLORIC   ACID;   HISTORY. 


CHAPTER   IX. 

HYDROCHLORIC  ACID. 

Formula,  H  Cl ;  specific  gravity,  air  =  1,  is  1.  2658,  H2  =.-2,  is  36.45  ; 
1  c.c.  at  0°  and  .76  m.  weighs  .0016442  gram. 

HYDROCHLORIC  acid  very  seldom  occurs  free  in  nature,  and 
then  only  in  the  exhalations  of  some  volcanoes  and  in  the  springs 
arising  from  their  craters ;  for  instance,  the  Eio  Vinagre,  arising  in 
the  Andes,  is  said  to  contain  .08  per  cent,  the  Paramo  de  Ruiz,  in 
New  Granada,  .8  per  cent  of  free  hydrochloric  acid. 

The  aqueous  solution  of  the  acid  was  first  prepared  by  Basil 
Valentine  in  the  fifteenth  century,  by  distillation  of  salt  (NaCl, 
sodium  chloride)  with  ordinary  green  vitriol  (FeS04,  ferrous  sul- 
phate) ;  it  was  subsequently  investigated  by  a  number  of  alchemists, 
but  the  pure  gas  was  not  obtained  until  1772,  when  Priestley  isolated 
pure  hydrochloric  acid.  The  old  name  was  spiritus  sails,  or  acidum 
muriaticum  (from  murias,  sea  salt),  and  at  the  present  day  the 
aqueous  acid  is  frequently  termed  muriatic  acid.  When  first  inves- 
tigated, hydrochloric  acid  was  supposed  to  contain  oxygen;  but 
Davy,  in  1810,  proved  that  it  was  composed  of  nothing  but  hydrogen 
and  chlorine. 

The  acid  is  best  prepared  by  treating  the  chloride  of  a  metal, 
preferably  sodium  chloride,  with  sulphuric  acid,  when  the  following 
reaction  takes  place  :  — 

2  Nad  +H2S04  =Na2S04  +2HC1 

Sodium  chloride + Hydrogen  sulphate  =  Sodium  sulphate + Hydrogen  chloride.29 

The  gas  can  be  collected  over  mercury,  or,  being  heavier  than 
air,  by  downward  displacement,  but  not  over  water,  as  it  is  ex- 
tremely soluble  in  that  substance. 

Hydrochloric  acid  is  a  colorless  gas  with  an  acid  odor ;  it  fumes 
in  the  air,  owing  to  its  power  of  condensing  moisture  from  the 
atmosphere  to  form  an  aqueous  solution  of  hydrochloric  acid ;  it 
cannot  be  breathed,  as  it  causes  violent  coughing;  it  is  neither 


HYDROCHLORIC   ACID;   PREPARATION.  67 

combustible,  nor  will  other  substances  burn  in  it.  )  The  stability  of 
union  of  hydrogen  and  chlorine  is  very  great ;  at  about  1800°  (high 
white  heat)  hydrochloric  acid  begins  slightly  to  decompose  into 
hydrogen  and  chlorine,  but,  as  we  have  become  aware  of  the  great 
tendency  which  hydrogen  has  to  unite  with  chlorine,  this  stability 
is  not  unexpected.  Hydrochloric  acid  is  very  soluble  in  water ;  at 
0°  one  volume  of  water  will  dissolve  505  volumes  of  hydrochloric 
acid  gas ;  the  solution  then  contains  43  per  cent  of  the  acid.30 

(When  chlorine  is  brought  in  contact  with  hydrogen,  in  the  dark, 
no  reaction  takes  place ;  if  the  mixture  of  the  two  gases  is  exposed 
to  the  sunlight,  or  if  a  lighted  taper  is  applied,  or  an  electric  spark 
allowed  to  pass  through  the  gases,  a  violent  explosion  takes  place 
and  hydrc^lilcTic_aci^  is  produced.  On  the  other  hand,  if  a  current 
of  electricity  is  passed  through  a  concentrated  solution  of  hydro- 
chloric acid,81  the  chlorine  will  separate  at  the  positive  pole,  the 
hydrogen  at  the  negative ;  this  resembles  the  decomposition  of 
water,  excepting  that  with  hydrochloric  acid  equal  volumes  of 
hydrogen  and  chlorine  are  produced,  while  in  the  case  of  water  two 
volumes  of  hydrogen  and  one  of  oxygen  result.  When  hydrogen 
and  chlorine  are  mixed  in  equal  volumes  and  then  exploded,  there 
is  no  change  in  volume,  but  the  mixture  of  gases  is  entirely  con- 
verted into  hydrochloric  acid;  furthermore,  sodium,  when  it  is 
placed  in  a  closed  volume  of  hydrochloric  acid,  will  absorb  the 
chlorine  (forming  NaCl,  sodium  chloride)  and  the  volume  of 
the  gas  will  be  diminished  by  one-half.82} 

We  have  now  proved  that  hydrochloric  acid  decomposes  into 
equal  volumes  of  hydrogen  and  chlorine,  and  that  equal  volumes  of 
hydrogen  and  chlorine  unite  to  form  hydrochloric  acid,  the  resulting 
volume  being  equal  to  the  sum  of  the  volumes  of  hydrogen  and 
chlorine  before  union.  The  relationship  between  the  volumes  of 
hydrogen  and  oxygen  which  are  capable  of  producing  water  without 
leaving  a  residue  of  either  gas  is  equally  simple ;  so,  indeed,  is  that 
between  the  volumes  of  any  gases  uniting  to  form  a  gaseous  com- 
pound, so  that  Gay  Lussac,  who  first  accurately  investigated  the 
-subject,  was  able  to  formulate  the  following  law :  — 

"  Two  gases  always  unite  in  such  a  way  that  their  volumes  bear 
I  a  simple  ratio  to  each  other,  and  the  volume  of  the  resulting  prod- 
I  uct,  if  it  is  a  gas,  also  bears  a  simple  relationship  to  the  volumes 
I  of  the  original  gases." 


68  KINETIC   GAS   THEORY. 

According  to  the  kinetic  theory  of  the  nature  of  gases,  now  uni- 
versally held,  these  substances  are  composed  of  particles  of  matter 
which  are  flying  about  in  right  lines,  until  they  impinge  on  the 
sides  of  the  vessel  in  which  the  gas  is  contained,  or  on  each  other, 
when,  being  perfectly  elastic,  they  rebound.  The  particles  of  the 
gas  are  separated  by  such  distances  that  their  own  volume  exer- 
cises no  influence  on  the  volume  of  the  gas  as  a  whole.  Now,  the 
weight  of  a  given  volume  of  gas  is  but  the  sum  of  the  weights  of  the 
individual  particles  making  up  that  gas ;  and  the  specific  gravity  of 
any  gas,  if  air  is  taken  as  unity,  is  equal  to  the  weight  of  a  given 
volume  of  that  gas  as  compared  with  that  of  the  same  volume  of 
air.  Investigation  has  shown  us  that  the  ratios  between  the  spe- 
cific gravities  of  elementary  gases,  and  those  between  their  com- 
bining weights,  bear  a  simple  relationship  to  each  other ;  thus,  the 
specific  gravity  of  hydrogen  in  round  numbers,  is  .07,  of  oxygen, 
1.12  ;  but  .07  is  to  1.12  as  1 : 16,  and  1  part  of  hydrogen  unites  with 
8  parts  of  oxygen  (1  with  ^)  to  form  water.  The  specific  gravity 
of  chlorine  is  2.46,  and  the  relationship  is :  .07  :  2.45 :  :  1  : 35.1 ; 
but  the  combining  weight  of  chlorine  is  35.45,  as  that  part  by 
weight  of  chlorine  unites  with  one  part  of  hydrogen.  If  we  use 
hydrogen  as  the  standard  instead  of  air,*  the  relationship  is  seen 
more  readily,  thus  :  — 

H  =  1.     Then  the 

combining  weight  of  oxygen   is    8.  (or  ^6-)  its  specific  gravity  16. 
"  «       "  chlorine  is  35.45  "         "  "       35.45 

"  "       «  nitrogen  is    4.66  (or  *£)  its "  "      14. 

From  these  numbers  it  will  be  seen  that  equal  volumes  of  gases 
bear  a  similar  relationship  by  weight  to  each  other,  as  do  the  indi- 
vidual particles  of  which  they  are  composed,  and,  therefore,  it  is 
reasonable  to  suppose  that  the  numbers  of  particles  in  the  gases 
themselves  must  bear  a  simple  ratio  to  each  other,  as  1 : 2,  1 : 3,  or 
1 : 1,  indeed  the  last  ratio  (1  : 1)  is  the  most  probable  one,  for,  by 
constructing  a  theory  that  in  equal  volumes  of  gases  there  are 
equal  numbers  of  particles,  we  can  most  readily  explain  the  sim- 
ple relationship  which  exists  between  the  combining  volumes  of 
gases.  This  was  the  conclusion  reached  by  Gay  Lussac  in  the 

*  The  ratio  between  the  weights  of  hydrogen  and  of  an  equal  volume  of  air, 
is  as  .069:1  or  as  1:14.4,  hence,  any  specific  gravity  with  air  as  unity  can 
be  converted  to  hydrogen  as  unity,  by  multiplying  by  14.4. 


RELATIONS  OF  GAS  VOLUMES. 


69 


course  of  his  investigations.  Thus,  as  equal  volumes  of  hydrogen 
and  chlorine  unite  to  form  hydrochloric  acid,  and  if  the  ratio  be- 
tween the  weights  of  equal  volumes  of  the  gases  is  the  same  as  that 
between  the  weights  of  the  individual  particles,  it  follows  that  the 
volumes  of  the  two  gases  contained  equal  numbers  of  these  parti- 
cles before  their  union.  Dalton,  however,  soon  pointed  out  a  defect 
in  this  reasoning.*  Let  us  suppose  we  have  a  volume  of  hydrogen 
containing  1000  atoms,  then,  according  to  the  theory,  an  equal  vol- 
ume of  chlorine  will  also  contain  1000  atoms,  the  two  unite,  thus 
forming  1000  molecules  of  hydrochloric  acid.  The  natural  result 
would  be, 


or :  —  1  volume  hydrogen      +      1  volume  chlorine      =  1  volume  hydrochloric  acid, 

for,  as  we  have  seen,  the  volume  of  the  molecule  exercises  no 
influence  on  the  volume  of  the  gas.  Nature,  however,  contradicts 
this  reasoning,  for  we  know  that  hydrogen  and  chlorine  unite  with- 
out change  of  volume ;  in  other  words,  1  volume  of  hydrogen  -f- 1 
volume  chlorine  =  2  volumes  hydrochloric  acid^  and  it  follows 
that  if  hydrogen  and  chlorine  have  equal  numbers  of  atoms  in 
equal  volumes,  then  hydrochloric  acid  must  have  but  one-half  the 
number  in  the  same  volume,  for, 


500 
HC1 

500 
HC1 

1  volume  hydrogen       +     1  volume  chlorine      =  2  volumes  hydrochloric  acid, 

*  The  example  cited  by  Dalton  in  order  to  refute  Gay  Lussac's  argument, 
was  nitric  oxide  and  not  hydrochloric  acid.  Hydrochloric  acid  could  not  have 
been  referred  to  by  Dalton,  as  he  supposed  that  the  compound  contained  oxy- 
gen. The  same  relationship  by  volume  exists  between  nitrogen,  oxygen,  and 
nitric  oxide  as  obtains  between  hydrogen,  chlorine,  and  hydrochloric  acid,  so 
that,  for  the  purpose  of  the  present  argument,  either  gas  will  do  equally  well. 

t  Let  the  pupil,  instead  of  using  the  expression  "  volume,"  substitute 
u  liter,"  and  the  whole  subject  will  appear  more  clear,  thus, 

1  liter  hydrogen  +  1  liter  chlorine  =  2  liters  hydrochloric  acid. 


70 


AYOGADRO'S   HYPOTHESIS. 


so  1000  molecules  of  hydrochloric  acid  must  occupy  twice  the 
volume  previously  taken  by  1000  atoms  of  hydrogen,  and  hence  in 
one  volume  of  hydrochloric  acid  there  must  be  but  500  molecules. 
It  was  left  for  an  Italian  physicist,  Arnadeo  Avogadro,  to  explain 
this  seeming  discrepancy  between  theory  and  fact.  Avogadro  sup- 
posed the  elementary  gases  to  be  composed  of  molecules  instead  of 
atoms.  As  a  usual  thing,  these  molecules  are  composed  of  two 
atoms,  so  that,  accepting  this  hypothesis,  we  would  arrive  at  the 
following  result :  — 


200 

OHC1 

1  volume  hydrogen       +    1  volume  chlorine      =    2  volumes  hydrochloric  acid. 

A  reaction  of  this  kind  would  then  consist  simply  of  an  inter- 
change of  the  atoms  composing  the  molecules,  so  that,  whereas  we 
previously  had  molecules  each  of  which  was  composed  of  atoms 
of  the  same  kind,  we  now  would  have  molecules  each  of  which  is 
composed  of  atoms  which  are  of  a  different  kind.  This  will  be 
more  apparent  if  we  write  the  equation  as  follows :  — 

H— H  H     H 

+         =       I       I 
Cl— Cl  Cl    Cl 

What  is  true,  then,  of  the  volume  is  true  of  the  individual  mole- 
cule, there  being  the  same  number  of  molecules  in  equal  volumes ; 
the  terms  volume  and  molecule  can  therefore  be  used  interchange- 
ably. 

Simple  as  Avogadro's  explanation  was,  it  was  not  generally  ac- 
cepted by  chemists,  chiefly  because  he  tried  to  apply  it  in  cases 
where  substances  which  never  were  in  a  gaseous  state  were  con- 
cerned, so  that  it  was  not  until  many  years  later  that  it  was  uni- 
versally adopted.  It  was  then  brought  into  prominence  and  is  now 
one  of  our  fundamental  hypotheses,  furnishing  to  us  the  best  means 
of  determining  the  molecular  weights  of  chemical  compounds  and 
elements. 

Let  us  suppose  that  we  have  a  volume  of  hydrogen  weighing 


AVOGADRO'S    HYPOTHESIS. 


71 


two  grams,  then  an  equal  volume  of  chlorine  must  weigh  70.9  grams, 
for  the  weights  of  equal  volumes  of  elementary  gases  are  to  each 
other  as  their  atomic  weights  and  1  :  35.45  : :  2 :  70.9. 


HC1 

72. 

HC1 

9  grams. 

1  volume  hydrogen  +  1  volume  chlorine    =    2  volumes  hydrochloric  acid. 

It  follows  that  a  volume  of  hydrochloric  acid  equal  to  that  of  the 

72  9 
hydrogen  taken  must  weigh  — ~  grams  or  36.45   grams,  and,  as  in 

equal  volumes  of  the  gases  there  are  equal  numbers  of  molecules, 
the  ratio  between  the  weights  of  the  individual  molecules  of  hy- 
drogen and  hydrochloric  acid  must  be  2  : 36.45,  and  therefore,  if  the 
molecule  of  hydrogen  is  two,  then  the  molecule  of  hydrochloric  acid 
is  36.45 ;  so  that  if  hydrogen  is  taken  as  a  standard  and  placed  at 
two,  the  specific  gravity  of  hydrochloric  acid  is  equal  to  its  molec- 
ular weight,,  and  we  shall  soon  see  that  this  rule  can  be  made 
general  as  follows  :  —  •***" 

If  hydrogen  be  placed  at  two,  then  the  molecular  weights  and 
specific  gravities  of  gases  are  the  same  number.* 

Hydrogen  and  oxygen  unite  to  form  water  in  the  proportion  of 
two  volumes  of  hydrogen  to  one  of  oxygen  and,  if  the  water  so 
formed  is  measured  in  the  form  of  vapor,  we  shall  find  that  two 
volumes  of  this  vapor  are  produced,  thus :  — 


H 

1  volume 

+ 

H 

1  volume 

+ 

0 

1  volume 

= 

2  volumes 

H20 

If,  according  to  Avogadro's  hypothesis,  oxygen  has  two  atoms  to 
the  molecule,  and  supposing  that  we  select  volumes  each  of  which 
contain  1000  molecules,  the  reaction  which  takes  place  will  be  as 
follows :  — 

*  In  dealing  with  the  specific  gravity  of  gases  it  is  not  necessary  to  deal 
with  exact  numbers.  Thus,  using  the  oxygen  standard  for  atomic  weights, 
hydrogen  is  1.008,  but,  for  all  practical  purposes',  the  decimal  can  be  neglected. 


72 


SPECIFIC    GRAVITIES    OF    GASES. 


or,  as  molecule  and  volume  can  be  used  interchangeably, 
H— H  H— H         H  H  H  H 

0—0 


V  V 

0  0 

1  Mol.  H  +  1  Mol.  H  +  1  Mol.  0  =  2  Mols.  H2  0. 


Now,  if  each  volume  of  hydrogen  weighs  two  grams,  then  the  same 
volume  of  oxygen  will  weigh  thirty-two  grams,  and  consequently  :  — 


2  grams 
H 

+ 

2  grams 
H 

+ 

32  grs. 
0 

= 

18  grs. 
H20 

+ 

18  grs. 
H20 

Therefore,  in  the  case  of  water  also,  if  hydrogen  is  placed  at 
two,  the  molecular  weight  and  the  specific  gravity  are  identical,  and 
similar  methods  of  reasoning,  backed  by  experiment,  have  shown 
the  same  to  be  true  in  regard  to  the  specific  gravities  of  all  gases.* 

The  whole  of  the  preceding  reasoning  can  be  summed  up  as 
follows  :  — 

If  in  equal  volumes  of  all  gases  there  are  equal  numbers  of  mol- 
ecules, then  the  weights  of  these  equal  volumes  must  bear  the  same 
relationship  toward  each  other  as  do  the  weights  of  the  individual 
molecules.  The  relative  weights  of  the  molecules  are  therefore  to 
be  ascertained  by  determining  the  relative  weights  of  equal  vol- 
umes of  gases ;  and,  for  obvious  reasons,  the  weight  of  a  molecule 
of  hydrogen  has  been  selected  as  the  standard  ;  as  the  molecule  of 
hydrogen  consists  of  two  atoms,  this  standard  for  measuring  molec- 
ular weights  is  placed  at  two.  If  the  specific  gravities  of  gases  have 
been  determined  with  air  as  unity,  then,  in  order  to  recalculate  them 
so  as  to  compare  them  with  hydrogen  as  two,  they  must  be  multi- 
plied by  28.8,  for :  - 

spec.  grav.  of  hydrogen  :  spec.  grav.  of  air  : :  2  :  28.8. 

The  value  of  the  discoveries  just  cited  as  an  assistance  in  determin- 
ing the  atomic  weights  of  elements  is  apparent.    By  a  determination 
*  For  apparent  exceptions  see  ammonium  chloride. 


SPECIFIC   GRAVITIES   OF   GASES.  73 

of  the  specific  gravity  of  a  gas,  we  ascertain  the  relative  .weight  of 
a  molecule  of  that  gas  as  compared  with  the  weight  of  a  molecule 
of  hydrogen  ;  as,  for  instance,  in  the  case  of  water,  the  molecular 
weight  cannot  be  more  or  less  than  eighteen,  and  in  this  eighteen 
parts  by  weight  of  water,  quantitative  analysis  shows  us  that  we 
have  sixteen  parts  by  weight  of  oxygen  and  two  of  hydrogen.  The 
atomic  weight  of  oxygen,  therefore,  cannot  be  more  than  sixteen, 
unless  we  wish  to  accept  the  existence  of  a  fraction  of  an  atom  of 
oxygen  in  a  molecule  of  water.  The  maximum  number  for  the 
atomic  weight  of  oxygen  is  consequently  fixed  by  experiment. 
That  it  is  not  some  fraction  of  sixteen  we  cannot  state  so  definitely, 
yet  all  evidence  points  against  this  assumption.  In  the  first  place, 
two  volumes  of  hydrogen  unite  with  one  of  oxygen  to  form  water ; 
the  presumption,  therefore,  is  that  water  has  the  formula  H2  0,  and 
hence  sixteen  would  be  the  atomic  weight  of  oxygen;  and  in  the 
second,  we  never  have  encountered  any  compound  of  oxygen  which 
can  be  obtained  in  a  gaseous  state,  and  whose  molecular  weight  we 
therefore  know,  which  contains  relatively  less  than  sixteen  parts  by 
weight  of  that  element.  The  magnitudes  at  present  in  use  for  our 
atomic  weights  are  the  results  of  reasoning  exactly  similar  to  that 
given  in  the  case  of  oxygen,  assisted  in  many  cases  by  deductions 
drawn  from  analogies  existing  between  a  number  of  elements  and 
by  other  less  important  methods  of  determining  molecular  weights, 
and  we  shall  subsequently  see  that  the  atomic  weights  at  present  in 
use  are  the  only  ones  by  means  of  which  the  elements,  when  arranged 
in  the  order  of  their  atomic  weights,  naturally  fall  into  series  and 
families  which  show  the  greatest  resemblance  to  one  another.  This 
existence  of  elements  as  molecules  is  used  to  explain  the  chemical 
activity  of  elements  in  the  nascent  state.  (See  page  51.) 

Chemical  equations,  expressing  the  changes  which  take  place 
when  simple  or  compound  substances  are  brought  in  contact,  indi- 
cate the  initial  bodies  and  the  final  result  by  formulae  based  upon 
our  atomic  weights,  taking  no  account  of  the  changes  of  energy ; 
they,  as  a  rule,  represent  only  the  main  course  of  a  reaction,  while 
other  minor  reactions  are  often  going  on  in  a  mixture  of  two  or 
more  substances.  Many  equations  are  true  only  for  certain  condi- 
tions of  temperature,  concentration,  etc.  A  chemical  equation  is 
simply  an  algebraic  expression  which  can  be  constructed  quite  inde- 
pendently of  chemical  facts,  and  when  so  constructed  is  entirely 


74  HYDROCHLORIC    ACID;    THKH  MO-CHEMISTRY. 

useless  if  not  pernicious  in  its  tendency.  We  must  always  bear  in 
mind  that  chemical  facts  and  experiments  are  infinitely  more  valu- 
able than  chemical  equations,  the  latter  being  useful  only  as  a  short 
method  of  expressing  those  changes  which  we  know  to  take  place. 

In  uniting  to  form  hydrochloric  acid,  hydrogen  and  chlorine 
give  220  K.  (  See  page  64.)  In  dissolving  in  water  an  additional 
173  K  is  evolved,  so  that  the  solution  of  hydrochloric  acid  in  water 
possesses  much  less  chemical  energy  than  does  the  gas.  We  should 
therefore  expect  hydrochloric  acid  gas  to  enter  into  a  number  of  re- 
actions where  the  solution  would  be  inert.  If  hydrochloric  acid 
gas  and  oxygen  are  passed  through  a  heated  tube  (see  page  59)  the 
following  reaction  takes  place  :  — 

0  =  H20  +  2  Cl 


while  if  chlorine  water  is  allowed  to  stand  in  the  sunlight  :  — 
2  Cl  +  H2  0  =  2  H  Cl  -f  0. 

This  contrast  is  explained  by  the  difference  in  energy  between 
gaseous  hydrochloric  acid  and  the  solution,  for  :  — 

H  +  Cl  =  H  Cl  gives  220  K, 
2H  +  2C1  =  2HC1  «  440  K, 
2H-f  0  =  H20  «  684  K, 

hence,  the  heat  of  formation  of  water  (the  measure  of  the  chemical 
energy  of  H  and  0)  is  greater  than  that  of  two  molecules  of  hydro- 
chloric acid  gas,  and  therefore  the  reaction  2  H  Cl  +  0  =  H2  0  + 
.2  Cl  would  be  accompanied  by  an  evolution  of  heat.  On  the  other 
;hand, 

H  +     Cl  =     H  Cl,  dissolved  in  water  gives  393  K, 
2  H  +  2  Cl  =  2  H  Cl         «          «      «         «     786  K, 

and  hence  the  heat  of  formation  of  dissolved  hydrochloric  acid  is 
greater  than  that  of  water.  As  a  result  the  reaction, 

2  C1  +  H20=2  HC1  +  0, 

in  the  presence  of  water  is  accompanied  by  an  evolution  of  heat.  Of 
course,  the  more  concentrated  a  solution  of  hydrochloric  acid  is,  the 
nearer  will  it  approach  the  condition  of  the  gas,  and  hence  the 
greater  will  be  its  reactiveness.  Similar  studies  with  other  bodies 
show  us  that  the  difference  in  energy  between  dissolved  substances 
and  those  undissolved  is  often  very  marked,  and  in  considering 


ACIDS;   NATURE   OF.  75 

whether  certain  chemical  reactions  will  take  place  we  should  take 
tliis  difference  into  account. 

The  compound  of  hydrogen  and  chlorine  is  called  an  acid 
because  it  has  certain  distinctive  properties  possessed  by  every 
substance  entitled  to  be  classed  as  such. 

An  acid  is  a  compound  containing  hydrogen  united  to  a  negative 
element  or  group  of  elements,  which  hydrogen  can  be  replaced  by  a 
metal  to  form  a  salt. 

No  definition  of  an  acid  is  entirely  satisfactory,  as  we  have 
a  number  of  substances  which  contain  hydrogen  replaceable  by  a 
metal,  as,  for  instance,  water  in  the  reaction  HOH  -}-  K  =  KOH  -(-  H, 
yet  we  scarcely  would  call  water  an  acid,  nor  KOH  a  salt,  although, 
essentially,  there  is  no  difference  between  this  reaction  and  the 
following  :  Zn  +  2  H  Cl  =  Zn  C12  +  2  H  ;  in  one  case,  it  is  the  neg- 
ative group  of  elements  OH,  in  the  other,  the  negative  element  Cl, 
which  is  united  to  hydrogen,  and  the  reactions  take  place  because 
K  or  Zn  has  a  greater  chemical  energy  when  brought  in  contact 
with  OH  or  Cl,  respectively,  than  has  hydrogen  ;  they  being  more 
metallic  in  their  nature  than  is  hydrogen,  and  hence  presenting  a 
greater  contrast  to  the  not-metal.  If  we  call  hydrochloric  acid  an 
acid,  and  water  not  one,  it  is  simply  because  expediency  shows  us 
that  it  is  well  to  classify  those  hydrogen  compounds,  the  hydrogen 
of  which  is  easily  replaced  by  a  large  number  of  metals  to  form 
salts,  under  the  head  of  acids;  an  indication  of  the  propriety  of  the 
name  being  that  the  substance  designated  as  "  acid  "  has  the  power 
of  turning  a  vegetable  dye  '(blue  litmus)  to  a  red  color.  It  is  evi- 
dent that  this  latter  distinction  is  purely  arbitrary,  and  unimportant 
as  regards  the  true  chemical  nature  of  acids.  Many  substances 
which  are  not  acids  will  turn  blue  litmus  to  a  red  color,  while,  on  the 
other  hand,  some  substances  decidedly  acid  *  have  no  effect  what- 
ever upon  litmus. 

In  forming  chlorides  we  can  employ  four  leading  methods,  three 
of  which  illustrate  general  characteristics  of  acids  ;  they  are  :  — 

1st.    By  direct  union  of  the  elements,  as,  for  example  :  — 


P    +3C1=PC13, 

Zn  +2  Cl  =  ZnCl2,  C     +4C1 


Organic  substances  acting  as  acids  (aceto-acetic  ether). 


76  CHLORIDES;   FORMATION   OF. 

This  method  leads  to  the  formation  of  chlorides  of  the  not- 
metals  as  well  as  those  of  the  metals,  the  process  being  analogous 
to  that  of  combustion  in  oxygen. 

2d.  By  the  action  of  hydrogen  chloride  on  a  metal,  by  which 
means  the  chloride  is  produced  and  hydrogen  liberated,  as  :  — 

Zn  +2HCl  =  ZnCl2  +  2  H, 
Fe  -f2  HCl  =  FeCl2  +  2  H, 
Mg  +  2  H  Cl  =  Mg  C12  +  2  H. 

This  reaction  takes  place  with  metals  only. 

3d.  By  the  action  of  hydrogen  chloride  on  the  oxides  of  metals, 
when  the  chloride  and  water  are  produced,  as  :  — 

ZnO  +2HCl  =  ZnCl2  +H20, 
CaO  +  2HCl  =  CaCi2  +H20, 
MgO  +2  HC1  =  MgCl2  +  H20. 

4th.  By  the  action  of  hydrogen  chloride  on  the  hydroxides  of  the 
metals,  as  :  — 

OH  HC1  Cl  HOH, 

Zn(  +  =Zn(  + 

XOH  HC1  XC1  HOH, 

Ca(OH)2+.2HCl  =CaCl2  +2H20, 

Mg(OH)2  +  2HCl  =MgCl2  +2H20. 

When  an  oxide  or  hydroxide  reacts  in  the  above  manner  it  is 
the  oxide  or  hydroxide  of  a  metal,  and  is  designated  as  a  base  ; 
while  the  chemical  process  of  forming  a  salt  by  addition  of  an  acid 
to  a  base  or  base  to  an  acid  is  called  neutralization  (the  acid  or  base 
is  neutralized).  This  term  "  base  "  is  one  dictated  by  expediency  ; 
and  when  we  speak  of  a  substance  as  basic  in  character,  we  mean  it 
can  form  an  oxide  or  a  hydroxide  which  will  neutralize  acids. 

The  reactions  under  3  and  4  are  general  to  all  acids  and  bases ; 
if  we  designate  any  metal  by  M,  any  acid  by  HX,  then  the  general 
rule  will  be  that :  — 

M20  +2HX=2MX  +  H20;  M  OH  +  HX=MX  +  H20; 
MO  +2HX=  MX2+  H20;  M  (OH)2  +  2HX=MX2+2H20; 
M203-f  6HX=2MX3  +  3H20;  M(OH)3-f  3HX=MX3+3H20. 


NEUTRALIZATION.  77 

Thus  :  - 

K2   0        +2HN03  =  2KN08  +    H20; 

Zn   0         +  2HBr    =ZnBr2  +    H20; 
A12  03       +  6HNO.  =  2Al(NO,),+3HaO; 

KOH         +    HN03  =  KN03  +    H20; 

Zn(OH)2+2HM)3  =  Zn(N03)2  +2H20; 

Al(OH),  +  3Htf08  =  Al(N08)8  +  3H20. 


The  metals  differ  among  each  other  in  their  power  of  replacing 
hydrogen  in  acids  to  form  salts  ;  some  replace  one  atom,  some  two, 
some  three,  to  form  one  molecule  of  the  salt;  but  whatever  the 
acid,  this  number  is  always  the  same  for  any  given  metal,  and  it 
can  be  ascertained  from  the  formula  of  the  chloride  of  the  metal 
(as  KC1,  ZnCl2,  A1C13)  ;  the  metal  will  replace  as  many  hydrogen 
atoms  in  any  acid  as  there  are  chlorine  atoms  in  the  formula  weight 
of  its  chloride,  and  the  hydroxide  will  contain  as  many  hydroxyl 
groups  (OH  )  as  there  are  chlorine  atoms  in  the  chloride.* 

The  reactions  under  2  take  place  between  a  number  of  metals 
and  acids,  but  the  applications  are  much  less  general  than  3  and  4, 
and  often,  indeed,  where  the  acid  contains  oxygen,  no  hydrogen  is 
evolved,  but  some  other  substance  is  produced  from  the  acid  by  the 
action  of  hydrogen  in  the  nascent  state. 

*  See  pages  30,  32,  and  43. 


78  BROMINE;  OCCURRENCE,  PREPARATION. 


CHAPTER   X. 

BROMINE   AND    HYDROBROMIC   ACID. 

*°. 

Symbol,  Br  ;  atomic  weight,  79.95  ;  specific  gravity  of  fluid,  3.187  at 

0°;  specific  gravity  of  gas,  below  900°,  tm-  =  l,  is  5.54,  H2=  2, 
is  159.5.  Formula,  HBr;  specific  gravity,  air  =  1,  is  2.81, 
H2  =  2,ts  80.95;  1  c.c.  a*  0°  a^  .76  m.  pressure  weighs  .00365 


IN  many  respects  bromine  resembles  chlorine  ;  indeed,  such 
modifications  in  chemical  characteristics  as  it  represents  are  due 
simply  to  its  larger  atomic  weight.  It  has  the  same  tendency  to 
unite  with  metals  to  form  salts,  and  hence  is  not  free  in  nature,  but 
is  always  found  combined  with  metals  in  the  form  of  bromides  ;  its 
compounds  generally  accompany  those  of  chlorine,  yet  they  are 
always  present  in  lesser  quantity,  bromine  being  one  of  the  rarer 
elements.  The  bromides  of  sodium  and  magnesium  are  found  in 
the  great  majority  of  salt  springs,  especially  in  those  of  Saratoga 
Springs,  in  the  Saginaw  Valley,  and  in  the  southeastern  portion  of 
Ohio,  where  the  bromide  of  potassium  also  occurs.  In  Europe  the 
brines  from  the  salt  works  of  Kreuznach  and  Strassfurth  are 
especially  rich  in  bromides.  Marine  fauna  and  flora  also  fre- 
quently contain  bromides.  The  brines  of  the  various  salt  works 
are  evaporated,  thus  crystallizing  the  sodium  chloride  for  the  man- 
ufacture of  table  salt,  there  remaining  a  not-crystallizable  brine 
(mother  liquor)  which  is  especially  rich  in  bromides.  In  this 
mother  liquor  Ballard  discovered  bromine  in  1826,  calling  it 
bromine  from  fip&fjLos,  a  stench. 

Bromine  is  prepared  from  its  compounds  in  a  manner  entirely 
analogous  to  the  method  used  in  isolating  chlorine.  A  bromide  is 
mixed  with  manganese  dioxide  and  sulphuric  acid  (see  pages  59,  60), 
when  the  following  reaction  takes  place  :  83  — 

H2S04-f  Mn  02  =  Mn  S04  +  Na2  S04  +  2  Br  +  2  H2  0. 


Hydrobromic  acid  and   manganese   dioxide  would   yield  bromine 


BROMINE;    PROPERTIES.  79 

just    as  hydrochloric  acid   and    manganese   dioxide  give  us   chlo- 
rine :  — 


4  H  Cl  +  Mn  02  =  Mn  C12  +  2  Cl  +  2  H2  0 

but  owing  to  the  difficulty  of  preparing  hydrobromic  acid,  the  latter 
substance  is  an  expensive  article,  so  that,  from  reasons  of  economy, 
this  method  is  not  available. 

Bromine  is  a  dark  brown  liquid,  almost  black  when  any  consid- 
erable thickness  is  observed;  it  melts  at  —  7°.  3  and  boils  at  63°.  05, 
a  point  considerably  below  the  boiling  point  of  water.  ,  The  specific 
gravity  of  the  liquid  at  0°  is  3.18.  'When  the  liquid  is  allowed  to 
stand  in  the  air  it  evaporates  rapidly,  even  at  ordinary  tem- 
peratures, yielding  reddish  brown  vapors  very  irritating  to  the 
mucous  membrane  of  the  eyes,  nose,  and  throat,  and  possessing  an 
odor  much  resembling  that  of  chlorine.  The  specific  gravity  of  the 
vapor,  air  being  1,  at  800°,  is  5.54,  giving  159.5  as  its  density,  H 
=  2.  This  shows  that  at  this  temperature  the  molecule  of  bromine 
is  composed  of  two  atoms  like  that  of  chlorine,  for  the  atomic 
weight  is  80,  so  that  160  would  be  the  molecular  weight  of  Br2  . 
About  1200°  the  specific  gravity  of  the  vapor  decreases  until  it 
.becomes  3.7,  showing  that  at  high  temperatures  some  of  the  Br2 
molecules  have  changed  to  individual  atoms. 

Bromine  is  soluble  in  water,  the  solution  having  a  brownish 
color  and  properties  similar  to  those  of  chlorine  water,  one  part  of 
bromine  at  15°  being  soluble  in  33  parts  of  water.  If  this  solution 
is  cooled  to  the  freezing  point  of  water,  crystals  of  a  compound  of 
bromine  with  water  of  crystallization  (  2  Br  +  10  H2  0  )  separate. 
(See  page  62.) 

The  solution  of  bromine  in  water  is  an  oxidizing  agent,  and  hence 
bleaches  just  as  chlorine  water  does  ;  this  property  is  due  to  the 
same  cause,  the  liberation  of  oxygen. 


The  formation  of  oxygen  becomes  apparent  if  a  tube  containing 
bromine  water  is  inverted  over  a  water  trough  and  exposed  to  the 
sunlight  ;  oxygen  separates  as  it  does  in  the  case  of  chlorine  water, 
although  not  with  such  great  rapidity.* 

*  Owing  to  the  greater  ease  with  which  bromine  is  handled,  it  is  more  fre- 
quently in  use  as  a  laboratory  oxidizing  agent  than  is  chlorine. 


80  HYDROBROMIC   ACID;   PROPERTIES. 

The  compounds  of  bromine  resemble  those  of  chlorine  in  every 
particular,  and  the  bromides  and  chlorides  of  the  same  metal  are 
isornorphous.* 

Bromine  does  not  unite  with  hydrogen  as  readily  as  does  chlo- 
rine, its  higher  atomic  weight  rendering  its  chemical  character  less 
negative,  and  hence  its  tendency  to  unite  with  metals  less  pro- 
nounced ;  as  a  consequence,  a  mixture  of  bromine  and  hydrogen  can 
be  allowed  to  stand  in  the  sunlight  for  any  length  of  time  without 
the  formation  of  hydrobromic  acid  ;  the  union  is  only  to  be  brought 
about  by  more  energetic  means,  such  as  the  electric  spark,  or  the 
passing  of  a  mixture  of  hydrogen  and  bromine  over  platinized 
asbestos. 

The  heat  of  formation  of  hydrobromic  acid  gas  f  is  only  121  K, 
while  that  of  gaseous  hydrochloric  acid  is  220  K,  so  that  we  should 
expect  hydrobromic  acid  to  be  more  easily  decomposed  than  is 
hydrochloric  acid.  The  consequences  of  this  instability  are  un- 
pleasantly apparent  in  the  difficulties  encountered  in  the  prepara- 
tion of  hydrobromic  acid. 

In  preparing  hydrochloric  acid  we  had  but  to  treat  sodium  chlo- 
ride with  sulphuric  acid,  as  follows  :  — 


and  a  similar  reaction  takes  place  when  a  bromide  is  substituted 
for  a  chloride  :  —  ^^  •C^x>JM(v>^> 

2  ]STa  Br  +  H2  S04  =  Naa  S04  +  2  H  Br  ; 

but  hydrobromic  acid  (being  so  much  less  stable  than  hydrochloric), 
owing  to  the  heat  of  the  reaction,  breaks  down  into  hydrogen  and 
bromine,  after  which  decomposition  the  nascent  hydrogen  attacks 
the  sulphuric  acid,  reducing  the  latter  to  form  sulphur  dioxide  and 
water  :  -  1.  2  H  Br  =  2  H  -f-  2  Br 

2.  H2  S04    +  2H  =  2H20  +  S02; 

and,  as  a  consequence,  the  hydrobromic  acid  produced  by  this 
means  is  contaminated  with  sulphur  dioxide. 

Preparation  of  hydrobromic  acid. 

In  order  to  prepare  hydrobromic  acid  for  laboratory  use,  advan- 
tage is  taken  of  the  instability  of  the  halogen  compounds  of  the 
not-metals. 

*  See  page  42  and  foot-note.  t  Using  gaseous  bromine. 


HYDROBROMIC   ACID  ;    PKOPERTIES.  81 

When  phosphorus  trichloride  is  added  to  water,  the  following 
change  takes  place  :  — 


/jCl  +  HjOH          /OH 
P  —  JC1  +  H;0  H  =  P  —  OH  +  3HC1. 
\C1  +  HJOH          \OH 


Phosphorus  trichloride  +  water  =  phosphorus  hydroxide  (phosphorous  acid) 
+  hydrochloric  acid. 

P  Cl,  +  3  H20  =  P  (OH  )3  +     3  H  Cl. 
The  same  with  phosphorus  tribromide  : 

P  Br3  +  3  H20  =  P  (OH  )8  -f-  3  H  Br.34 

In  performing  this  operation  it  is  not  necessary  to  employ  the 
finished  tribromide  of  phosphorus,  for  a  mixture  of  phosphorus, 
bromine,  and  water  will  answer  the  same  purpose. 

Hydrobromic  acid  is  a  colorless  gas,  with  an  acid  odor  resem- 
bling that  of  hydrochloric  acid  ;  it  fumes  strongly  in  the  air,  owing 
to  the  absorption  of  moisture  ;  it  is  extremely  soluble  in  water,  one 
part  of  that  substance  absorbing  as  much  as  82  per  cent  of  hydro- 
bromic  acid,  and,  as  a  consequence,  the  gas  cannot  be  collected  over 
water.  If  a  quantity  is  desired,  it  must  either  be  separated  by  col- 
lecting over  mercury,  or  by  the  displacement  of  air  in  some  vessel, 
for,  as  the  specific  gravity  of  hydrobromic  acid  is  2.79,  it  can  be 
poured  downward  in  the  atmosphere. 

As  has  already  been  stated,  hydrobromic  acid  is  much  less 
stable  than  is  hydrochloric  acid*  ;  and  therefore,  on  adding  chlorine 
to  the  former,  bromine  is  separated  and  hydrochloric  acid  is 


In  its  chemical  behavior,  hydrobromic  acid  is  like  hydrochloric 
acid.     When  brought  in  contact  with  bases,  it  'neutralizes  them  to 
form  salts.          Na  QH  +     H  Br  =  Na  Br     +  HOH 
KOH    +    HBr  =  KBr       +  HOH 
CaO      +2HBr  =  CaBr2     +H20. 


*  Hydrobromic  acid  is  partly  broken  down  into  its  constituents  if  it  is  ex- 
posed to  the  sunlight;  the  same  is  true,  in  a  much  smaller  degree,  of  hydro- 
chloric acid,  if  it  is  in  concentrated  solution  in  the  presence  of  oxygen.  (A. 
PJchardson;  Journal  Chem.  Soc.  ;  51,  801.) 


\ 

82  BROMIDES  ;   FORMATION. 

The  methods  of  formation  of  the  bromides  are  analogous  to 
those  of  the  chlorides.  They  are  :  — 

1.  Direct  union  of  the  element  in  question  with  bromine,  this 
applying  to  metal  or  not-metal. 

2.  The  addition  of  a  metal   to   hydrobromic   acid,   when  the 
bromide  is  formed  and  hydrogen  liberated. 

3.  The  action  of  hydrobromic  acid  on  the  oxides  or  hydroxides 
of  the-metals,  when  the  bromides  and  water  are  produced. 

Hydrobromic  acid  is  formed  of  one  volume  of  hydrogen  and  one 
volume  of  bromine  vapor,  united  to  form  2  volumes  of  hydrobromic 
acid ;  the  same  considerations  advanced  under  chlorine,  show  that 
the  bromine  molecules,  provided  the  temperature  be  not  too  high, 
consist  of  two  atoms  to  the  molecule.  (See  pages  69,  70,  71.) 


IODINE;    OCCURRENCE,    HISTORY.  83 


CHAPTER  XL 

IODINE   AND    HYDROIODIC    ACID. 

Symbol,  I ;  atomic  weight,  126.85 ;  specific  gravity,  4.948  ;  specific 
weight  of  vapor,  air  =  1,  is  8.84  (below  600°),  H2  =  2,  is  255.21. 
Formula,  HI ;  specific  weight,  air  =  1,  is  4.44,  H2  =  2,  is  127.9. 
1  c.c.  of  the  gas,  0°  and  .76m.,  weighs  .005767  gram. 

(THIS  last  member  of  the  halogen  family  is  not  found,  as  such,  in 
nature ;  in  that  way  it  resembles  fluorine,  chlorine  and  bromine ;  its 
compounds,  although  they  occur  in  company  with  those  of  chlorine 
and  bromine  in  almost  all  deposits  in  which  the  halides  of  the 
metals  are  found,  are  not  by  any  means  present  in  such  large  quan- 
tities. The  element,  in  combination,  occurs  in  sea  water,  both  as  the 
iodide  of  sodium  and  of  magnesium,  although  the  quantities  of  these 
salts  are  so  small  that  their  presence  cannot  be  proved,  excepting 
by  some  special  means.  Sea  plants,  such  as  the  algae,  as  well  as 
representatives  of  the  animal  kingdom  (sea  sponges,  crabs,  oysters, 
etc.),  can  assimilate  and  concentrate  traces  of  iodides  so  that,  on 
drying  and  burning,  iodine  can  easily  be  proved  to  be  present  in 
their  ashes.  The  iodides  occur  in  salt  springs,  in  deposits  of  rock 
salt,  in  a  number  of  mineral  springs,  as  at  Kreuznach  and  Reichen- 
hall,  in  river  water,  and  also  in  some  water  plants  growing  in 
flowing  fresh  water. 

The  element  was  discovered  in  1811  by  Courtois,  a  saltpetre 
manufacturer  in  Paris,  who  found  its  compounds  in  the  mother- 
liquors,  left  after  extracting  the  ashes  of  sea  plants  and  crystalliz- 
ing the  less  soluble  portions.  The  name  iodine  is  taken  from 
iwSr/s,  violet.  The  weed  which  is  washed  up  by  the  spring  storms 
on  the  islands  on  the  western  coast  of  Scotland  or  of  Ireland,  or  on 
the  coast  of  Normandy,  or  that  which  grows  upon  the  rocks,  is  dried 
and  burned,  the  fused  mass  remaining  as  the  ash  is  brought  into 
the  market  under  the  Scotch  name  of  kelp,  or  Normanic,  varec. 
The  amount  of  iodine  contained  in  this  ash  is  very  small,  and  is  ex- 
tracted from  the  last  remaining  mother-liquors,  obtained  by  crystal- 


84  IODINE;  PREPARATION,  PROPERTIES. 

lizing  the  aqueous  extract  of  kelp.  (The  method  of  preparation  of 
iodine  is  identical  with  that  of  chlorine  or  bromine ; 35  the  iodide  is 
treated  with  sulphuric  acid  and  manganese  dioxide : 


rf 


Mn  O2  +  2  KI  +  2H2  S04  =  MnS04  +  K2  S04  +  2  H2  0  +  2 1. 

In  order  to  purify  the  iodine  it  is  sublimed,  the  iodine  being  heated 
in  retorts  and  collected  as  crystals  in  cold  chambers. 

v  The  element  is  almost  black,  grayish  solid,  with  a  lustre  closely 
resembling  that  of  the  metals ;  when  pure  and  fused  it  is  entirely 
black ;  in  thin  plates  it  is  translucent  with  a  brownish-red  color) 
Its  specific  weight  is  4.94 ;  it  melts  at  113°— 115°  and  boils  at  about 
200° ;  *  when  heated  in  a  vacuum  it  does  not  melt,  but  vaporizes 
without  fusion.  The  vapor  of  iodine  has  a  beautiful  violet  color  f 
and  a  specific  gravity,  air  being  1,  of  8.84,  below  600°,  giving  a 
density,  H2  =  2,  of  255.21,  showing  that  below  this  temperature 
iodine  has  a  molecule  consisting  of  two  atoms,  I2 ;  but  if  the  heat  is 
gradually  increased,  the  specific  gravity  of  the  vapor  diminishes,  so 
that  at  1570°  it  is  only  (air  =  1)  5.67,  or  (H2  =  2)  163.69,  indicat- 
ing that  the  vapor  of  iodine  at  bright  red  heat  consists  of  a  mixture 
of  the  individual  atoms,  and  the  molecules  I2 ,  dissociation  beginning 
above  600°.$  Iodine  is  but  very  little  soluble  in  water,  about  7000 
parts  of  water  dissolving  one  part  of  iodine ;  water  containing  salts 
in  solution  has  a  greater  solvent  action;  some  solutions,  such  as 
those  of  potassium  iodide  and  of  hydroiodic  acid,  have  the  power  of 
dissolving  large  quantities  of  iodine ;  the  element  is  also  extremely 
soluble  in  substances  such  as  alcohol,  ether,  carbon  bisulphide,  or 
chloroform. 

i  Chemically,  iodine  resembles  chlorine  or  bromine  ;  it  unites  with 
sulphur,  phosphorus,  and  other  not-metals  with  which  the  other 
lialogens  form  compounds ;  in  combination  with  the  metals  it  forms 
iodides ;  and  if  hydrogen,  mixed  with  iodine  vapors,  is  passed  over 

*  Stas;  Gesetze  der  Chemischen  Proportionen ;  141. 

t  Observe  the  same  by  throwing  some  iodine  on  a  hot  stove  or  into  a  hot 
porcelain  crucible.  When  the  vapor  of  iodine  is  saturated,  it  is  pure  blue  in 
color. 

\  At  1600°-1700°,  H.  Biltz  and  Y.  Meyer  find  the  vapor  density  of  iodine  to 
correspond  with  a  molecular  weight  which  indicates  that  at  the  temperature 
of  the  experiment  the  molecules  I2  have  completely  broken  down  into  the 
individual  atoms.  (Berichte  derDeutsch.  Chem.  Gesell.  22;  726.) 


HYDROIODIC   ACID  ;   PREPARATION.  85 

spongy  platinum  which  is  heated,  some  hyclroiodic  acid  is  formed, 
although  hydroiodic  acid  is  an  endothermic  compound,  and  there- 
fore possesses  more  energy  than  the  individual  elements  of  which 
it  is  composed.  } 

When  the  preparation  of  hydrobromic  acid  was  mentioned,  we 
saw  that  the  method  available  for  hydrochloric  acid  was  not  feasi- 
ble, owing  to  the  relative  instability  of  hydrobromic  acid  ;  the  acid 
breaking  down  into  hydrogen  and  bromine,  while  the  hydrogen  re- 
duces the  sulphuric  acid  to  sulphurous  acid  ;  as  a  consequence  we 
are  compelled  to  use  a  reaction  which  depends  on  the  decomposition 
of  the  bromide  of  phosphorus  by  water.  We  could  not,  therefore, 
expect  to  obtain  hydroiodic  acid  by  the  action  of  sulphuric  acid  on 
sodium  iodide,  for  hydroiodic  acid  decomposes  even  more  readily 
than  does  hydrobromic  ;  in  this  case  the  reduction  of  sulphuric  acid 
takes  place  so  energetically  that  sulphuretted  hydrogen  (H2  S  )  is 
produced.  If  we  write  the  equations  representing  the  action  of 
hydrobromic  and  hydroiodic  acid  upon  sulphuric  acid,  this  relative 
instability  at  once  becomes  apparent  :  — 

H2  S04  +  2  H  Br  =  2  H2  0  +  S0a       +  2  Br, 
H2S04  +  8HI     =H2S     +  4H20+8I. 

for  in  the  latter  case  the  decomposition"  is  much  more  energetic  and 
far-reaching.* 

;  In  the  preparation  of  hydroiodic  acid,  therefore,  (we  are  com- 
pelled to  resort  to  a  round-about  method  similar  to  that  employed 
in  the  production  of  hydrobromic  acid  ;  the  reaction  which  is  used 
depends  on  the  instability  of  the  iodide  of  phosphorus  in  the  pres- 
cnceofwate, 


*  This  reaction  illustrates  quite  well  the  influence  exerted  by  the  mass  of 
the  chemical  reagents  used.  If  a  large  excess  of  sulphuric  acid  and  but  little 
hydroiodic  are  present,  a  considerable  quantity  of  sulphur  dioxide  (SO2)  is  pro- 
duced, while  generally  more  or  less  sulphur  is  deposited.  These  changes  are 
represented  by  the  equations  :  — 

H2S04  +  6HI  =  4H2Q  +  S      +61. 

H2  SO4  +  2  HI  =  2  H.,  O  +  SO2  +  21. 

Under  ordinary  conditions,  in  a  test  tube,  all  of  these  changes  take  place  at  the 
same  time,  so  that  the  equation  given  in  the  text  only  represents  one  of  the 
various  changes  going  on;  a  variation  of  temperature  or  mass  of  the  reagents 
can  alter  the  proportions  of  hydrogen  sulphide,  sulphur,  or  sulphur  dioxide 
produced, 


86  HYDROIODIC    ACID;   PROPERTIES. 

Of   course,   a  mixture    of   iodine,  water,    and   red   phosphorus 

answers  the  purpose.36   (Hydroiodic  acid  can  be  collected  in  empty 

"jars  by  displacement  of  air,  or  it  can  be  collected  over  mercury  ;  its 

extreme  solubility  in  water  renders  the  filling  of  vessels  with  the 

gas  impossible  where  water  is  present.) 

Hydroiodic  acid  is  a  colorless  gas,  with  the  acid  odor  of  hydro- 
chloric or  hydrobromic  acid ;  it  fumes  in  the  air,  owing  to  absorption 
of  moisture/  Its  specific  weight  (air  being  1),  is  4.44  ;  H2  being  2,  it 
is  127.9 ;  the  molecular  weight  of  HI  is  127.85,  the  slight  discrep- 
ancy between  the  observed  specific  gravity  and  the  molecular  weight 
being  due,  undoubtedly,  to  the  difficulties  encountered  in  obtaining 
pure  hydroiodic  acid.  Hydroiodic  acid  is  quite  easily  condensed  to 
a  liquid ;  at  —  17°. 8  c.  its  vapor  tension  is  but  two  atmospheres ; 
it  solidifies  at  —  51°  c.  On  heating,  hydroiodic  acid  is  easily  de- 
composed, the  change  into  hydrogen  and  iodine  beginning  at  100°, 
"  and  being  complete  at  700°.*  (Chlorine  decomposes  hydroiodic 
acid  with  almost  explosive  violence,  forming  hydrochloric  acid  and 
iodine.fj  * 

Hydroiodic  acid  is  composed  of  equal  volumes  of  hydrogen  and 
iodine ;  two  volumes  of  hydroiodic  acid  yield  one  of  hydrogen  and 
one  of  iodine,  provided  the  temperature  be  kept  sufficiently  high  to 
vaporize  the  iodine  formed.  The  gas  is  extremely  soluble  in  water ; 
the  solution,  when  saturated  at  12°,  contains  57.7  percent  of  hydro- 
iodic acid ;  it  must  be  kept  well  stoppered  and  in  the  dark ;  for, 
when  exposed  to  the  air,  a  separation  of  iodine  takes  place,  owing 
to  the  formation  of  water,  as  follows :  — 


(  Hydroiodic  acid  has  a  strongly  acid  reaction ;  it  turns  blue  lit- 
mus solution  red,  and  in  all  respects  resembles  hydrochloric  and 
hydrobromic  acids,  the  methods  of  formation  of  the  iodides  being 
exactly  like  those  of  bromides  and  chlorides.  The  heat  of  forma- 
tion of  gaseous  hydroiodic  acid  from  hydrogen  and  solid  iodine  is 

*^Bodenstein  ;  Berichte  d.  Deutsch.  Chem.  Gesell. ;  26,  2603. 

t  The  iodine  first  separates  as  the  element,  and  this  is  then  converted  into 
the  trichloride  of  iodine,  a  yellow  solid.  As  a  consequence,  the  violet  vapors 
which  at  first  appear  when  chlorine  is  added  to  hydroiodic  acid  subsequently 
disappear,  while  the  trichloride  of  iodine  settles  on  the  sides  of  the  vessel  in 
which  the  reaction  takes  place. 


HALOGENS  ;   COMPARATIVE   TABLE   OF. 


87 


—  61  K ;  from  gaseous  iodine  and  hydrogen,  almost  nil ;  the  heat  of 
solution  is  192  K,  so  that  the  heat  of  formation  of  hydroiodic  acid 
in  water  is  131  K.  The  gaseous  acid  is  therefore  an  endo thermic 
compound,  possessing  a  tendency  to  decompose  into  its  constituent 
parts ;  the  solution  possesses  much  less  chemical  energy,  and  hence 
a  greater  stability.  In  the  following  table  a  comparison  of  the 
properties  of  the  halogens  has  been  undertaken,  their  differences 
with  increasing  atomic  weight  thus  being  made  more  apparent :  — 


THE  HALOGENS. 


F. 

Cl. 

Br. 

I. 

Pale  yellow 

Greenish  yel- 

Dark 

Black 

gas. 

low  gas. 

brown 

solid. 

liquid. 

Density  of  liquid. 



1.33 

3.18 

4.97 

Density        <  Air  =  1 

1.26 

2.46 

5.54 

8.84 

of  vapor.    (   H2  =  2 

36.28 

70.84 

159.5 

255.21 

Molecule   of  gaseous 

element. 

F2 

C12 

Br2 

I* 

The  density  of  iodine  vapor  at  447°  is  8.8 ;  at  1570°  the  density 
is  5.67,  which  yields  163.69  as  the  density,  H2  =  2.*  Dissociation 
is  therefore  far  advanced  at  this  temperature,  so  that  at  a  higher 
one  the  molecule  and  atom  of  iodine  nearly  correspond. 


HF.           HC1.                HBr.                   HI. 

Stability. 

J===— 

Heat    of    for- 

mation. 



H,  Cl  =  220  K. 

H,Br=121K. 

H,  I  =  -  61  K. 

Ditto,  plus 

water. 



393  K. 

320  K. 

131  K. 

*  A  determination  by  V.  Meyer  (Berichte  d.  Deutsch,  Chem.  Gesell. ;  13, 
1310),  taken  at  1437°  gives  a  density  of  4.76  for  the  vapor  of  iodine;  this, 
H2  =  2,  would  give  a  specific  gravity  of  137.4,  or  a  molecular  weight  which  very 
nearly  corresponds  to  the  atomic  weight  of  iodine.  See  also  foot-note  page 
84.  This  refers  to  the  last  work  on  the  subject. 


88       HALOGENS;  COMPARATIVE  TABLE  OF. 

Chlorine  replaces  bromine  when  brought  in  contact  with  bromides,  and 
iodine  when  brought  in  contact  with  iodides. 
Bromine  replaces  iodine  in  the  iodides. 
Fluorine  liberates  oxygen  from  water,  even  in  the  dark. 

2F  +  H20=2HF  +  0. 

Chlorine  liberates  oxygen  from  wTater  in  the  sunlight. 
2C1  +  H2O=2HC1  +  O. 

Bromine  liberates  oxygen  from  water  in  the  sunlight  more  slowly  than 
chlorine. 

Iodine  does  not  decompose  water. 


THE   OXYGEN   FAMILY.  89 


CHAPTER   XII. 

THE   OXYGEN   FAMILY. 

THREE  elements  —  sulphur,  selenium,  and  tellurium  —  show  a 
relationship  toward  oxygen  similar  to  that  displayed  by  chlorine, 
bromine,  and  iodine  toward  fluorine.  With  increasing  atomic 
weights  we  have  similar  changes  in  the  physical  properties  of  the 
elements,  illustrating  the  decreasing  not-metallic  characteristics  of 
the  family ;  oxygen  is  a  colorless  gas,  sulphur  a  yellow  solid,  sele- 
nium a  dark  red  solid  (in  one  of  its  allotropic  forms),  while  tellurium 
is  an  element  having  entirely  the  appearance  of  a  metal.  All  of 
the  elements  of  this  family  form  hydrogen  compounds,  which,  with 
the  exception  of  water,  are  colorless  gases  at  ordinary  temperatures ; 
in  this  they  resemble  the  halogens,  for  in  that  family  the  hydrogen 
compound  of  the  element  with  the  smallest  atomic  weight  (fluorine) 
is  liquid  at  ordinary  temperatures.  The  compounds  with  hydrogen, 
which  are  produced  by  the  members  of  the  oxygen  family,  have 
one  atom  of  the  element  united  to  two  of  hydrogen ;  the  atoms  of 
the  elements  under  consideration,  therefore,  display  twice  as  great 
a  power  of  retaining  hydrogen  atoms  in  close  proximity  as  the 
atoms  of  the  halogens  do  ;  the  formulae  of  these  hydrogen  compounds 
are  H20,  H2S,  H2Se,  H2Te;  their  stability  and  heat  of  formation 
diminish  with  increasing  atomic  weight,  accompanying  the  decreas- 
ing not-metallic  properties  of  the  elements.  This  deportment  is 
exactly  parallel  to  the  similar  changes  with  the  members  of  the  hal- 
ogen family,  so  that  hydrogen  selenide  decomposes  at  about  the 
same  temperature  as  hydroiodic  acid  (150°),  while  hydrogen  tellu- 
ride  is  not  known  in  a  pure  state,  and  decomposes  even  at  ordinary 
temperatures  if  placed  in  contact  with  the  air.  The  hydrogen 
compounds  of  this  family  are  by  no  means  as  acid  in  their  proper- 
ties as  are  those  of  the  halogens,*  —  indeed,  water  does  not  redden 
litmus,  while  hydrogen  sulphide,  selenide,  and  telluride  do  not  turn 

*  The  acid  properties  of  the  hydrogen  compounds  of  the  not-metals  dimin- 
ish with  an  increase  in  the  number  of  hydrogen  atoms  in  the  molecule. 


90  THE   OXYGEN   FAMILY. 

blue  litmus  into  a  pronounced  red.  The  hydrogen  compounds  of 
sulphur,  selenium,  and  tellurium  are  much  less  soluble  in  water 
than  hydrochloric,  hydrobromic,  and  hydroiodic  acids  are. 

In  comparing  hydrofluoric  acid,  water,  and  ammonia  (H3N),  we 
see  that  hydrofluoric  acid,  as  its  name  implies,  is  an  acid ;  water 
acts  like  an  acid  only  when  it  is  brought  in  contact  with  the  most 
pronounced  metals,  such  as  sodium  or  potassium  (sometimes  it  can 
act  as  a  base),  while  ammonia  (H8K)  is  basic  in  its  character.  This 
gradation  of  properties  is  possibly  due  to  the  successively  increasing 
predominance  of  the  positive  element  in  the  three  substances,  for  it 
is  obvious  that  the  three  hydrogen  atoms  in  a  molecule  of  ammonia 
will  have  a  much  greater  effect  in  determining  the  character  of  the 
compound  than  the  one  present  in  a  molecule  of  hydrofluoric  acid ; 
hence  H3ISf  as  an  entire  compound  partakes  of  the  nature  of  a  metal 
and  is  positive.  As  the  negative  element  in  the  hydrogen  com- 
pounds increases  in  atomic  weight,  its  larger  mass  may  cause  its 
character  to  become  more  prominent ;  hence  H3  Sb  is  not  basic,  an- 
timony having  next  to  the  highest  atomic  weight  in  the  family  of 
which  nitrogen  is  a  member. 

The  study  of  the  hydrogen  compounds  of  the  elements  of  the 
oxygen  family  is  less  completely  finished  than  is  that  of  the  corre- 
sponding compounds  of  the  halogens,  because  hydrogen  selenide  and 
hydrogen  telluride  have  rarely  been  prepared  in  a  pure  state. 

The  elements  of  this  family,  with  the  possible  exception  of  tel- 
lurium, exist  in  two  allotropic  forms.  They  form  compounds  with 
oxygen  which  are  anhydrides  of  acids ;  the  study  of  these  will  be 
deferred  until  the  oxygen  compounds  of  the  not-metals  are  consid- 
ered. The  compounds  of  sulphur,  selenium,  and  tellurium  with 
the  metals  have  formulae  corresponding  to  those  of  the  oxides, 
thus  :  — 

Na20  Na2S  Na2Se  Na2Te 

Sodium  oxide ;  Sodium  sulphide ;  Sodium  selenide;  Sodium  telluride. 

FeO  FeS  FeSe  FeTe 

Ferrous  oxide ;  Ferrous  sulphide ;  Ferrous  selenide ;  Ferrous  telluride. 

Oxygen  has  already  been  discussed,  so  that  the  study  of  this 
family  will  be  continued  with  sulphur. 


SULPHUR;  OCCURRENCE,  HISTORY.  91 


CHAPTER  XIII. 

SULPHUR. 

Symbol,  S ;  atomic  weight,  32.06 ;  specific  gravity  of  solid,  2.045 ; 
specific  gravity  of  vapor  (above  1000°),  air  =  1,  is  2.2,  H2  =  2, 
is  63.36. 

SULPHUR  is  found,  often  in  company  with  gypsum  or  limestone, 
in  volcanic  regions  such  as  those  of  Southern  Italy  and  of  Sicily. 
In  Europe  the  largest  quantities  come  from  the  provinces  of  Gir- 
genti  and  Catania  in  Sicily ;  the  crater  of  the  volcano  Purace  in 
South  America,  with  a  surface  of  about  1200  yards  square,  is  cov- 
ered with  a  layer  of  sulphur  more  than  a  yard  in  thickness ;  while 
in  our  own  country  considerable  quantities  of  sulphur  are  found  in 
California.  The  element,  when  combined,  is  chiefly  found  in  the 
form  of  sulphides,  the  most  important  of  which  are  given  in  the 
following  table  :  — 

Iron  sulphides :  —  Iron  pyrites  and  markasite,  Fe  S2 . 

Lead  sulphide :— Galenite  (galena),  PbS. 

Zinc  Sulphide:  —  Zinc  blende,  ZnS. 

Copper  and  iron  sulphide:  —  Chalcopyrite  (CuFe)  S2. 

Sulphur,  united  with  oxygen  and  a  metal,  occurs  in  the  sulphates 
of  calcium,  CaS04,  barium,  BaS04,  magnesium,  MgS04,  etc. ;  the 
most  important  of  these  is  gypsum,  Ca  S04  +  2  H2  0.  Many  organic 
compounds  (such  as  mustard  oils)  contain  sulphur. 

Sulphur  has  been  known  from  very  ancient  times,  having  been 
used  as  a  medicine  by  the  Greeks  and  Romans.  The  alchemists 
considered  it  an  essential  portion  of  all  combustible  substances, 
while  during  the  period  in  which  credit  was  given  to  the  phlogiston 
theory,  it  was  looked  upon  as  a  compound  of  sulphuric  acid  with 
phlogiston.  After  our  present  chemical  theories  were  initiated  by 
Lavoisier's  studies  on  oxidation,  sulphur  was  classed  as  an  element. 

Of  course,  as  sulphur  is  found  uncombined,  any  method  for  pre- 
paring the  element  from  its  compounds  is  of  no  essential  value  in 
the  laboratory.  Sulphur  might  be  furnished,  just  as  is  chlorine, 


92  SULPHUR;    PREPARATION. 

bromine,  or  iodine,  by  treating  its  hydrogen  compound  with  some 
oxidizing  agent ;  for  instance,  by  mixing  a  sulphide  with  manganese 
dioxide  and  adding  sulphuric  acid  (a  process  similar  to  that  used  in 
the  preparation  of  the  halogens)  ;  but  such  a  method  would  be  of 
no  practical  value,  and  would  only  be  of  use  because  it  shows  the 
similarity  of  action  between  the  hydrogen  compounds  of  the  ele- 
ments of  the  family  under  discussion  and  those  of  chlorine,  bromine, 
and  iodine.  An  interesting  method  for  the  preparation  of  sulphur, 
which  is  important  because  it  explains  the  occurrence  of  the  ele- 
ment in  the  neighborhood  of  volcanoes,  is  by  the  action  of  sulphur 
dioxide  on  sulphuretted  hydrogen.37  The  reaction  takes  place  as 
follows:—  SQ2  +  2H2S  =  3S  +  2H20* 

Sulphur  dioxide  +  Sulphuretted  hydrogen  =  Sulphur  +  Water. 
Both  sulphur  dioxide  and  sulphuretted  hydrogen  occur  in  the  gases 
exhaled  from  volcanoes. 

If  iron  pyrites  (FeS2),  is  heated,  the  following  reaction  takes 
place  :  3  Fe  S2  =  Fe3  S4  -J-  2  S,t  so  that,  probably,  by  reason  of  this 
fact  there  exists  another  source  of  natural  sulphur.  The  heat 
necessary  to  decompose  iron  pyrites  might  have  been  furnished 
either  directly  by  subterranean  fires  or  by  some  process  of  oxida- 
tion going  on  throughout  the  entire  mass  of  sulphide. 

In  order  to  prepare  commercial  sulphur  from  the  rock  in  which 
the  element  occurs,  the  latter  is  broken  into  small  pieces  and  placed 
in  heaps  upon  the  side  of  a  hill ;  the  heaps  are  then  covered  with 
earth,  and  the  sulphur  lighted ;  the  heat  from  the  burning  portion 
is  sufficient  to  melt  the  remainder,  which  is  allowed  to  run  into  a 
receiver.  This  method,  in  vogue  in  the  sulphur  regions  of  Sicily, 
necessarily  involves  considerable  waste,  so  that  a  number  of  im- 
provements, a  description  of  which  belongs  in  a  more  extended  and 
technical  work,  have  been  introduced.  The  crude  sulphur  is  puri- 
fied by  distillation.  The  sulphur  of  commerce  occurs  in  two  forms, 
in  sticks  or  lumps,  and  as  a  yellow  powder  called  flowers  of  sulphur. 
The  latter  is  formed  by  rapidly  cooling  the  sulphur  vapors  upon  the 
walls  of  the  receiver  during  the  process  of  distillation. 

*  Compare  this  reaction  with  the  one  for  the  preparation  of  chlorine  by 
means  of  an  oxidizing  agent. 

t  Compare  this  reaction  with  the  one  for  the  preparation  of  oxygen  by 
heating  manganese  dioxide :  — 

3  Mil  O2  =  Mn3  O4  +  2  O. 


SULPHUR;  PROPERTIES.  93 

Sulphur  in  its  ordinary  form  is  a  solid  of  a  light  yellow  color. 
When  heated  it  melts  at  114°,  forming  an  amber  colored  liquid ;  on 
increasing  the  temperature  to  150°  the  molten  sulphur  darkens  in 
color  and  becomes  viscid,  so  that  at  200°  it  is  nearly  black,  and  so 
thick  that  it  can  no  longer  be  poured ;  at  about  340°  it  once  more 
becomes  fluid,  but  remains  of  a  dark  color.  Its  boiling  point  lies  at 
440°  centigrade,  at  which  temperature  it  changes  to  a  dark  brown 
vapor.38  The  specific  gravity  of  this  vapor,  air  being  1,  at  about  470°, 
is  7.8 ;  but  upon  increasing  the  temperature  this  gradually  dimin- 
ishes until  1000°  is  reached,  when  the  density  is  2.3.  At  the  latter 
temperature  the  molecule  of  sulphur  resembles  that  of  oxygen, 
hydrogen,  and  chlorine,  for  it  consists  of  two  atoms.* 

Sulphur  is  found  in  two  allotropic  forms ;  one  soluble  in  carbon 
bisulphide,  the  other  insoluble.  The  soluble  variety  can  be  par- 
tially changed  into  the  insoluble  one  by  melting  or  by  exposure 
to  the  sunlight.  Both  forms  occur  in  the  sulphur  of  commerce, 
flowers  of  sulphur  being  especially  rich  in  the  insoluble  variety. 

Sulphur  can  exist  in  four  crystalline  forms,  only  two  of  which 
are  easily  obtainable  and  of  importance. f  When  de- 
posited from  its  solutions  in  carbon  bisulphide,  it 
crystallizes  in  octahedra  belonging  to  the  rhombic 
system  (Fig.  7),  but  if  sulphur  is  melted  and  allowed 
to  cool  slowly,  it  forms  long  prismatic  needles  (mono- 
clinic  system),  as  in  Fig.  S.39  The  latter  form  changes 
into  the  former  on  standing ;  the  needles  becoming 
opaque,  and,  while  apparently  retaining  their  crystal- 
line form,  finally  consist  only 
of  an  aggregation  of  rhombic 
crystals.  During  this  process 
Fig-  8-  heat  is  evolved,  and,  therefore, 

*  It  is  generally  considered  that  the  molecule  of  sulphur  consists  of  six 
atoms  at  temperatures  just  above  the  boiling  point;  but  the  recent  investiga- 
tions of  Biltz  seem  to  show  that  sulphur  has  no  definite  vapor  density  below 
1000°.  Molecules  of  greater  complexity  than  S2  may  exist;  these  gradually, 
with  increasing  temperature,  decompose  into  simpler  ones.  The  specific  gravity 
of  7.8  indicates  (hydrogen  being  2)  a  specific  gravity  of  224.6,  which  corre- 
sponds to  a  molecule  S7 .  Compare  H.  Biltz ;  Zeitschrif  t  f  iir  Phys.  Chemie ;  3, 228. 

t  A  substance  which  exists  in  two  crystalline  forms  is  said  to  be  dimor- 
phous; one  existing  in  three,  trimorphous;  one  existing  in  many,  polymor- 
phous. For  a  description  of  the  four  crystalline  modifications  of  sulphur,  see 
W.  Muthmann;  Zeitschrif  t  fur  Krystallographie;  17,  336. 


94  SULPHUR;    CHEMICAL   BEHAVIOR    OF. 

the  rhombic,  stable  variety  of  sulphur  possesses  less  energy  than 
the  other. 

Sulphur,  when  heated  to  just  below  its  boiling  point  and  subse- 
quently rapidly  cooled  by  pouring  into  cold  water,  forms  a  plastic 
mass  somewhat  resembling  india  rubber  ;  this  gradually  becomes 
hard  and  changes  to  yellow  sulphur.  The  soft  sulphur  is  dark  in 
color,  probably  by  reason  of  impurities  ;  it  contains  both  soluble 
and  insoluble  sulphur  j  sulphur  crystals  are  sometimes  entirely 
soluble.* 

In  its  chemical  behavior,  in  many  respects,  sulphur  closely  re- 
sembles oxygen;  the  sulphides,  in  formula,  are  parallel  to  the 
oxides,  as  the  following  examples  demonstrate  :  — 

OXIDES.  SULPHIDES. 

Na20,  Na2S, 

K20,  K2S, 

Fe  0,  Fe  S, 

Fe304,  Fe3S4, 

-^  2  ^35  -*2  ^s, 

P205,  P2S5. 

Sulphur  combines  with  the  majority  of  the  elements  ;  and  the  sul- 
phides can  be  prepared  by  direct  union  of  the  elements,  just  as  the 
oxides  can.  As  further  examples  of  the  resemblance  between  sul- 
phur and  oxygen,  the  formulae  of  the  sulphhydrates,  which  corre- 
spond to  those  of  the  hydroxides,  are  most  interesting. 

HYDKOXIDES.  SULPHHYDRATES. 

NaOH,  NaSH, 

KOH,  KSH. 

The  sulphides  and  sulphhydrates  of  the  metals  resemble  the  bases  ; 
with  acids  they  form  salts  and  sulphuretted  hydrogen,  just  as  the 
bases  form  salts  and  water,  thus  : 


(NaSH+     HC1= 
HC1= 


*  Plastic  sulphur  is  often  formed  when  sulphur  is  separated  from  its  com- 
pounds by  chemical  means,  as,  for  instance,  in  the  oxidation  of  many  of  the 
sulphides  of  metals  with  nitric  acid. 

t  Many  sulphides  are  not  attacked  by  acids  under  ordinary  circumstances, 


SULPHIDES   AS   ACIDIC   ANHYDRIDES.  95 

A  number  of  the  sulphides  of  the  not-metals  resemble  the  anhy- 
drides of  acids,  for  they  form  salts  with  the  sulphides  of  the  metals  ; 
in  such  salts  the  oxygen  has  been  replaced  by  sulphur,  for 
example  :  — 

CS2  +Na2S  =  Na2CS3 

Carbon  disulphide  +  Sodium  sulphide  =  Sodium  thiocarbonate.* 

CO  2  -j-Na20  =  Na2C03 

Carbon  dioxide       +  Sodium  oxide       =  Sodium  carbonate. 

In  these  reactions  carbon  disulphide  and  sodium  sulphide  bear  the 
same  relationship  to  each  other  as  sodium  oxide  and  carbon  dioxide 
do.  Sulphur  also  resembles  chlorine,  bromine,  or  iodine,  for  the 
sulphides  can  be  prepared  in  the  same  way  as  the  halides  are  ;  for 
example,  when  phosphorus  and  sulphur  are  heated  together,  a  sul- 
phide of  phosphorus  is  formed,  just  as  is  the  case  if  phosphorus  is 
heated  in  chlorine,  for  in  the  latter  event  a  chloride  is  produced. 
Iron-filings,  heated  with  sulphur,  will  form  the  sulphide  of  iron, 
just  as  the  same  metal,  heated  with  chlorine,  will  produce  the 
chloride.  In  addition  to  the  resemblances  which  have  been  men- 
tioned, the  fact  can  be  cited  that  many  of  the  sulphides  of  the  not- 
metals  are  decomposed  by  water,  forming  an  acid  and  hydrogen 
sulphide,  just  as  the  chlorides  of  the  same  elements  are  decomposed, 
forming  an  acid  and  hydrogen  chloride,  for  example  :  — 

1.  P2  S3  +  6  H20  =  2  H3P03  +  3  H2S, 

2.  P2S5  +  8  H20  =  2  H3P04  +  5  H2S. 

1.  PC13  +  3H2O=     H3P03  +  3HC1. 

2.  PC15  +  4H20=     H3P04+5HC1. 


The  acid  formed  in  the  reactions  which  are  numbered  1  is  phosphor- 
ous acid,  in  the  reactions  which  are  numbered  2  is  phosphoric  acid. 
The  tendency  to  form  oxygen  compounds  of  the  not-metals,  in  most 
cases,  is  greater  than  is  the  one  to  form  compounds  containing  other 
elements  which  are  similar  to  oxygen. 

and  the  same  may  be  said  of  a  number  of  oxides.  Representatives  of  this 
class  are  found  as  minerals.  Whenever  the  sulphides  are  dissolved  by  acids, 
they  act  like  bases. 

*  The  term  thio  is  frequently  used  in  place  of  sulpho;  thus  sodium  sulpho- 
carbonate  is  more  often  termed  sodium  tMocarbonate  ;  See  chapter  xxxix. 


96  HYDKOGEN   SULPHIDE;    OCCURRENCE ;    HISTORY. 


CHAPTER  XIV. 

HYDROGEN   SULPHIDE. 

Formula,  H2S;  specific  gravity,  air  =  1,  is  1.17,  H2  =  2,  is  33.7 ;  1 
c.c.  of  gas  weighs  .00152  grams. 

HYDROGEN  sulphide  (sulphuretted  hydrogen)  occurs,  mixed 
with  other  gases  and  vapors,  in  some  volcanic  exhalations,  and  is 
also  occasionally  present  in  coal  and  in  other  mines.  As  it  is  one 
of  the  products  of  the  decay  of  various  animal  and  vegetable  sub- 
stances, it  is  necessarily  frequently  found  in  the  atmosphere  and  in 
water.  Many  sea  plants,  when  exposed  to  the  sun's  rays,  give  off 
sulphuretted  hydrogen  ;  and  when  organic  substances,  which  contain 
sulphur  (for  example,  bituminous  coal),  are  heated  without  access 
to  the  air,*  sulphuretted  hydrogen  is  found  in  the  gases  which  are 
given  off.  Many  mineral  waters  contain  large  quantities  of  sul- 
phuretted hydrogen ;  these  waters  come  from  sulphur  springs  ;  the 
gas  can  be  detected  by  its  peculiar  odor,  which  resembles  that  of 
rotten  eggs. 

Sulphuretted  hydrogen  has  certainly  been  known  since  the 
sixteenth  Or  seventeenth  century ;  at  a  later  time  it  was  more  accu- 
rately studied  by  Scheele,  who  considered  it  to  be  composed  of 
heat,  sulphur,  and  phlogiston.  After  Lavoisier's  time  its  true  nature 
was  explained. 

1.  Preparation  of  sulphuretted  hydrogen  by  direct  union  of  the 
elements. 

Sulphuretted  hydrogen  can  be  prepared,  with  some  difficulty,  by 
passing  hydrogen  through  molten  sulphur.  When  we  recall  the 
fact  that  hydrogen  unites  with  explosive  violence  with  either  oxy- 
gen or  chlorine ;  the  diminished  energy  with  which  hydrogen  and 
sulphur  or  hydrogen  and  bromine  combine  becomes  apparent. 

2.  Preparation  of  sulphuretted  hydrogen  for  laboratory  use. 

*  Called  dry  distillation. 


HYDROGEN   SULPHIDE;   PREPARATION.  97 

Sulphuretted  hydrogen  is  prepared  for  laboratory  use  by  the 
action  of  some  acid  on  a  sulphide  ;  for  instance  :  — 

Fe  S  +  H2S04  =  Fe  S04  +  H2  S, 
FeS+2HCl  =  FeCl2  +'H2S, 
H2S04  =  ZnS04  +  H2S. 


The  oxide  or  the  chloride  would  act  in  exactly  the  same  way  :  — 

FeO     +  H2S04  =  FeS04  +    H20, 
Fe  C12  +  H2  S04  =  Fe  S04  +  2  H  Cl, 

ZnO     +H,S04  =  ZnS04+     H20, 
Zn  C12  +  H2  S04  =  Zn  S04  +  2  H  Cl, 

so  that  there  is  no  essential  difference  between  the  action  of  a  base 
and  that  of  other  similarly  constructed  compounds  in  the  presence 
of  sulphuric  acid  ;  indeed,  were  sulphides  and  chlorides,  and  not 
oxides,  the  most  frequent  chemical  compounds,  or  were  hydrochloric 
acid  and  sulphuretted  hydrogen,  and  not  water,  produced  in  many 
reactions,  the  term  base  would  never  have  been  used  to  designate 
the  oxide.40 

Sulphuretted  hydrogen  is  a  colorless  gas,  with  an  intensely 
disagreeable  odor;  it  liquefies  at  a  temperature  of  11°,  with  a  pres- 
sure of  fifteen  atmospheres.  It  boils  at  ordinary  pressures  at  —  61°.  8 
and  becomes  solid  at  —  85°  ;  it  is  tolerably  soluble  in  water,  for  at 
ordinary  temperatures  one  volume  of  water  dissolves  two  volumes 
of  the  gas.  The  solution  slightly  reddens  litmus,  but  the  red  color 
disappears  on  exposure  to  the  atmosphere.  Sulphuretted  hydrogen 
is  very  poisonous  ;  when  inhaled  in  small  quantities  it  causes  head- 
ache, loss  of  appetite,  dizziness,  and  inflammation  of  the  eyelids, 
while  persons  who  have  been  poisoned  by  sulphuretted  hydrogen 
often  have  fainting  spells,  at  intervals,  during  some  weeks.  Death 
may  be  caused  by  -gfa  of  a  volume  of  sulphuretted  hydrogen  in  the 
atmosphere. 

When  a  stream  of  the  gas  is  lighted,  it  burns  to  form  water  and 
sulphur  dioxide  ;  a  mixture  of  two  volumes  of  sulphuretted  hydro- 
gen and  three  of  oxygen  is  highly  explosive  ;  where  an  insufficient 
supply  of  oxygen  is  present  during  the  combustion  of  the  gas, 
sulphur  is  deposited.  Both  chlorine  and  bromine  decompose 
sulphuretted  hydrogen  in  a  manner  similar  to  the  decomposition  of 


98  HYDROGEN    SULPHIDE;    THERMOCHEMISTRY. 

the  same  substance  by  oxygen,  forming  hydrobromic  or  hydrochloric 
acid :  — 

H2S  +     0   =     H20  +  S, 

H2  S  +  2  Cl  =  2  H  Cl  +  S, 

H2S-f  2Br  =  2HBr+S. 

In  each  of  these  three  cases  the  reason  for  the  reaction  is  found  in 
the  excess  of  the  heat  of  formation  of  water,  hydrochloric  acid,  and 
hydrobromic  acid  over  that  of  sulphuretted  hydrogen.  Iodine  does 
not  decompose  hydrogen  sulphide  when  no  water  is  present, 
because,  in  the  production  of  hydroiodic  acid,  heat  is  absorbed  ;  but 
if  the  reaction  takes  place  in  contact  with  water,  the  heat  of  solution 
of  hydroiodic  acid  is  sufficient  to  cause  the  following  change  to  take 

place:  H2S+2I=2HI  +  S. 

The  foregoing  changes  will  be  understood  if  the  heats  of  forma- 
tion of  these  various  compounds  are  placed  in  a  column  so  as  to 
render  comparison  easier :  — 

H20  =  684K, 
2HC1  =  440K, 
2  H  Br  =  242  K, 
2  HI     =  -  122  K,  while 
2  HI  dissolved  in  water  =  262  K. 

Now,  the  heat  of  formation  of  sulphuretted  hydrogen  is  but  27  K  ; 
the  heat  of  solution  of  the  same  in  water  is  46  K ;  therefore,  the  heat 
of  formation  of  hydrogen  sulphide  in  water  is  73  K,  a  number  much 
smaller  than  any  of  those  given  above.* 

These  relationships  become  apparent  if  we  write  the  complete 
reactions,  including  the  thermal  values  of  the  various  parts,  thus  :  — 

*  In  each  of  the  cases  in  which  hydrogen  sulphide  is  decomposed  by  a 
halogen,  the  reaction,  H2S  +  2  X  =  2  HX  +  S,  takes  place  (X  representing 
either  Cl,  Br,  or  I).  From  the  table  given  above,  it  follows  that  the  system, 
H2  S  +  2  X  possesses  more  energy  than  2  HX  +  S ;  heat  is  therefore  evolved 
when  chlorine,  bromine,  or  iodine  acts  on  hydrogen  sulphide,  and  hence  the 
reaction  takes  place.  In  each  of  these  cases  the  energy  which  must  be  added 
is  that  which  is  necessary  to  decompose  H2  S  into  2  H  +  S;  for  it  is  obvious  that 
hydrogen  sulphide  must  be  broken  down  into  hydrogen  and  sulphur  before  the 
rearrangement  with  the  new,  and  more  stable  bodies  (the  halhydric  acids)  can 
take  place;  the  energy  which  is  given  off  is  from  the  reaction  2  H  +2  X  =  2  HX; 
in  each  case  the  latter  must  exceed  the  former  if  a  change  is  to  take  place. 


HYDROGEN    SULPHIDE  ;    DECOMPOSITION.  99 

1.  H2S[27K]  -f  2C1  =  2H,C1[440K]-H2JS[27K]  +  S  = 
413  K. 

2.  H2  S  aq.  [73  K]  +  2 1  =  2  H,  I  aq.  [262  K]  -  H2,  S  aq.  [73  K] 
+  S  =  189  K. 

3.  H2  S  [27  KJ  +  0  =  H2,0  [684  K]  -  H2,  S  [27 K]  +  S  = 
657  K. 

The  term  aq.  used  above  in  equation  (2)  means  that  the  react- 
ing substances  are  dissolved  in  a  quantity  of  water  so  great  that  a 
further  addition  would  have  no  effect  on  the  thermal  value  of  the 
equation.  The  commas  placed  between  the  two  portions  of  a 
chemical  formula  indicate  that  the  substance  is  to  be  formed  from 
the  elements  in  question. 

Hydrogen  sulphide  is  readily  decomposed  into  its  elements,  just 
as  hydrobromic  and  hydroiodic  acids  are,  so  that  the  former  cannot 
be  prepared  in  the  presence  of  concentrated  sulphuric  acid  any  more 
than  the  two  latter  can,  and  it  follows  that  sulphuretted  hydrogen 
cannot  be  dried  by  being  passed  through  sulphuric  acid.  A  hot 
wire  placed  in  the  gas  will  dissociate  it,41  but  no  change  in  the 
volume  takes  place,  because  the  volume  of  the  sulphur  separated  is 
minimal  when  compared  with  the  volume  of  the  gas  as  a  whole, 
and,  therefore,  the  hydrogen  occupies  the  same  space  as  that 
previously  taken  by  sulphuretted  hydrogen :  — 

H  H 

x  S  =      +  sulphur. 

H  H 

1  volume  hydrogen  sulphide  =  1  volume  hydrogen. 

In  place  of  each  molecule  of  hydrogen  sulphide,  therefore,  we 
have  formed  one  molecule  of  hydrogen,  or,  in  place  of  each  volume 
of  hydrogen  sulphide,  an  equal  volume  of  hydrogen,  so  that  this 
fact  proves,  provided  Avogadro's  hypothesis  is  correct,  that  in  a 
molecule  of  sulphuretted  hydrogen  there  are  contained  at  least  two 
atoms  of  hydrogen. 

We  assume  that  there  is  but  one  atom  of  sulphur,  because  of  the 
analogy  between  water  and  sulphuretted  hydrogen,  and,  also,  because 
while  Ho  S  contains  two  parts  by  weight  of  hydrogen  and  thirty- 
two  of  sulphur,  in  no  compound,  the  molecular  weight  of  which  has 


100  HYDROGEN   SULPHIDE;    REACTIONS    OF. 

been  determined  and  which. contains  sulphur,  is  that  element  found 
with  a  proportional  weight  smaller  than  thirty-two. 

Hydrogen  sulphide  resembles  the  acids,  because,  in  cases  where 
it  reacts  with  the  oxide  or  hydroxide  of  a  metal,  the  corresponding 
sulphide  or  sulphydrate  is  formed  :  — 

CaO      +  H2S      =CaS     +  H20, 
Na  OH  +  HSH    =  NaSH  +  HOH. 

The  reactions  of  the  same  bases  with  hydrochloric  acid  are  as 

follows :  — 

CaO      +2HCl  =  CaCl2   +  H20, 
NaOH-fHCl     =  NaCl    +  HOH. 

Sulphuretted  hydrogen  contains  two  atoms  of  hydrogen,  which 
can  be  replaced  by  metals,  so  that  we  must  distinguish  two  series  of 
compounds,  in  the  first  of  which  one  of  these  atoms  is  substituted 
as  in  Na  SH ;  in  the  second,  two,  as  in  Ca  S.  We  shall  subsequently 
see  that  the  same  rule  holds  good  with  all  acids  containing  two 
replaceable  hydrogen  atoms. 

The  sulphides  of  a  large  number  of  metals  are  insoluble  in  dilute 
acids,  so  that  when  sulphuretted  hydrogen  is  passed  into  a  solution 
containing  a  salt  of  one  or  more  of  these  metals,  the  corresponding 
sulphide  is  precipitated :  — 

Cu  S04       +  H2  S  =  Cu  S  +  H2  S04, 
Pb  (N03)2  +  H2S  =  PbS+2H  N03, 
CdCl2        +  H2S  =  CdS+2HCl. 

The  sulphides  of  some  metals  are  soluble  in  dilute  acids,  but 
insoluble  in  water  or  in  alkalis ;  these  will  not  be  precipitated  unless 
provision  is  made  to  neutralize  the  acid  formed.  This  can  be  accom- 
plished either  by  passing  sulphuretted  hydrogen  into  a  solution  ren- 
dered alkaline  by  the  addition  of  ammonium  or  sodium  hydroxide 
solutions,  or,  better,  by  adding  a  soluble  sulphide. 

Na2  S  +  Zn  S04  =  Zn  S  +  Na2  S04. 

Lastly,  the  sulphides  of  a  third  class  of  metals  are  soluble  both 
in  acid  or  alkaline  water;  these  will,  of  course,  not  be  precipitated 
by  sulphuretted  hydrogen.  The  action  of  sulphuretted  hydrogen 
on  the  salts  of  the  metals  is  a  ready  means  of  detecting  the  presence 


HYDROGEN   PERSULPHIDE.  101 

of  the  metals  in  a  solution,  and  many  of  the  processes1  of '  quali'tairVe' 
analysis  are  founded  upon  these  reactions,  so  that  a  further  discus- 
sion belongs  to  that  branch  of  applied  chemistry. 

In  addition  to  hydrogen  sulphide  there  exists  another  compound 
of  hydrogen  and  sulphur  called  hydrogen  persulphide.  This  latter 
has  the  formula  H2  S2 ;  *  a  yellow,  oily  fluid  with  a  penetrating  odor 
and  corrosive  action.  It  resembles  its  prototype,  hydrogen  dioxide, 
(H202),  in  the  fact  that  it  is  stable  in  the  presence  of  dilute  acids, 
and  in  the  ease  with  which  it  decomposes  into  sulphur  and  sulphur- 
etted hydrogen,  just  as  hydrogen  dioxide  does  into  water  and  oxygen. 

*  This  may  be  the  formula  of  the  compound,  although  this  is  not  definitely 
settled.  There  may  exist  more  than  one  persulphide  of  hydrogen,  or  the  reason 
of  varying  amounts  of  sulphur  which  are  found  in  the  persulphide  may  be  due 
to  sulphur  dissolved  by  that  substance. 


102 '•'  '   SELENIUM;  OCCURRENCE,  HISTORY. 


CHAPTER   XV. 

SELENIUM   AND   HYDROGEN   SELENIDE. 

Symbol,  Se  ;  atomic  weight,  79.1 ;  specific  gravity,  4.5  ;  specific  grav- 
ity of  vapor,  above  1400°,  air  =  1,  is  5.7  ;  H2  =  2,  is  164  (calcu- 
lated for  Se2  158.2  ).  formula,  H2  Se  ;  specific  gravity,  air  =  1,  is 
2.8,H2=2,^80.6. 

SELENIUM  occurs  in  selenides  just  as  sulphur  is  found  in  sul- 
phides; in  quantity,  however,  it  is  found  much  more  sparingly. 
Free  selenium  is  only  very  rarely  found,  and  then  in  some  volcanic 
regions.  The  chief  selenides  are  those  of  lead  ( Pb  Se)  and  of  iron 
(FeSe),  while  the  selenide  of  silver  (Ag2Se)  also  occurs.  Lately, 
considerable  quantities  of  selenium  combined  with  bismuth  have 
been  found  in  some  parts  of  South  America. 

Selenium  was  discovered  by  Berzelius,  in  1817,  in  the  lead  cham- 
bers used  in  the  manufacture  of  sulphuric  acid.  He  at  first  con- 
fused the  element  with  tellurium,  but  subsequently  proved  it  to  be 
a  hitherto  unknown  element,  which  he  called  selenium  from  o-cXyv-r), 
moon,  because  the  name  of  the  other  element  was  derived  from 
tellus,  earth. 

The  occurrence  of  selenium  in  the  sulphuric  acid  chambers,  and 
in  the  flues  of  furnaces  in  which  sulphides  are  roasted,  is  due  to  the 
presence  of  selenium  in  the  ores  (such  as  iron  pyrites,  copper  pyrites, 
or  zinc  blende).  The  selenium  is  burned  to  selenium  dioxide,  and 
mechanically  carried  into  the  flues  and  chambers  when  the  sul- 
phides are  roasted.  Selenium  dioxide  is  easily  reduced  to  selenium 
by  means  of  reducing  agents  such  as  sulphur  dioxide,*  so  that  the 
dust  of  the  chambers  contains  selenium  chiefly  in  the  form  of  the 
element.  The  isolation  of  selenium  is  quite  a  complicated  process, 
and  a  description  of  the  methods  must  be  left  to  a  larger  work. 

Selenium  exists  in  two  allotropic  forms,  one  soluble,  the  other 

*  Reducing  agents  are  such  substances  as  are  capable  of  removing  oxygen 
or  the  equivalent  of  oxygen  from  chemical  compounds,  or  they  are  such  sub- 
stances as  can  add  hydrogen  to  elements  or  compounds. 


HYDROGEN    SELENIDE.  103 

insoluble  in  carbon  bisulphide ;  in  that  way  it  resembles  sulphur. 
When  selenium  is  separated  from  its  compounds  by  chemical  means 
it  is,  when  moist,  a  crimson  powder ;  but  it  is  dark  red  and  soluble 
in  carbon  bisulphide  when  dry.  By  heating  the  element  above  80° 
it  becomes  iron  gray  in  color.  Selenium  melts  at  217°,  and  boils  at 
665°,  and  then,  when  cooled  suddenly,  is  insoluble  in  carbon  bisul- 
phide. The  selenium  of  commerce  is  formed  by  casting  melted 
selenium  into  sticks.  It  is  almost  metallic  in  its  appearance,  and 
black  in  color.  Selenium  can  exist  in  more  than  one  crystalline 
form,  in  that  way  resembling  sulphur. 

Chemically,  the  properties  of  selenium  are  closely  akin  to  those 
of  sulphur.  It  burns  in  the  air,  forming  selenium  dioxide,  just  as 
the  latter  forms  sulphur  dioxide.  The  selenides  and  selenium 
compounds,  in  general,  have  formulae  exactly  like  those  of  the  cor- 
responding sulphur  compounds.  The  element  is  of  little  importance 
excepting  in  a  comparative  study  of  the  elements. 

Hydrogen  selenide  is  the  complete  analogon  of  hydrogen  sul- 
phide. It  can  be  prepared  with  difficulty  by  the  direct  union  of  the 
elements,  being  obtained  by  passing  hydrogen  over  selenium  heated 
to  its  boiling  point ;  however,  unless  great  care  is  taken  to  regulate 
the  temperature,  the  heat  will  decompose  the  hydrogen  selenide  so 
formed.  The  gas  can  also  be  prepared  by  adding  an  acid  to  the 
selenide  of  iron  :  — 

Fe  Se  +  2  H  Cl  =  Fe  C12  +  H2  Se, 

in  a  manner  analogous  to  the  preparation  of  hydrogen  sulphide. 

Hydrogen  selenide  is  a  colorless  gas,  with  a  most  penetrating 
odor,  somewhat  resembling  that  of  sulphuretted  hydrogen ;  it  is. 
extremely  poisonous.*  The  gas,  upon  being  heated,  begins  to.  "de- 
compose at  a  temperature  of  150°,  but  is  only  completely  dissociated 
at  a  considerably  higher  point.  It  burns  even  more  readily  than 
does  sulphuretted  hydrogen,  and,  of  course,  is  decomposed  by  chlo- 
rine or  bromine,  or  by  iodine  in  the  presence  of  water.  It  is  more 
soluble  in  water  than  is  sulphuretted  hydrogen,  and  has  the  same 
action  upon  soluble  salts  of  the  metals  as  the  latter. 

*  Care  must  be  taken  in  working  with  the  gas,  for  its  odor  clings  to  the 
clothes  for  many  days. 


104  TELLURIUM   AND   HYDROGEN    TELLURIDE. 


CHAPTER   XVI. 

TELLURIUM  AND   HYDROGEN    TELLURIDE  ;  COMPARATIVE 
TABLE   OP   THE   ELEMENTS   OF  THE   OXYGEN   FAMILY. 

Symbol,  Te ;  atomic  weight,  125  j  *  specific  gravity,  6.25 ;  specific 
gravity  of  vapor,  air  =  l,  is  9,  H2=2,  is  259  (above  1400°). 
Formula,  H2  Te. 

TELLURIUM  resembles  both  selenium  and  sulphur;  it  occurs  as  tel- 
luride  of  silver,  gold,  lead,  and  also  as  tellurium.  It  was  discovered 
in  1782,  and  identified  as  an  element  in  1798.  It  is  very  rare  and 
of  comparatively  little  importance.  Its  preparation  from  its  ores  is 
a  complicated  process.  It  is  a  silver  white,  metallic  appearing  ele- 
ment, \vhich  melts  at  about  500°,  and  boils  at  a  high  temperature, 
forming  an  orange-colored  vapor.  The  element  is  with  difficulty 
obtained  free  from  selenium.  Like  sulphur  and  selenium,  it  exists 
both  as  amorphous  and  crystalline  tellurium. 

Hydrogen  telluride  was  discovered  by  Davy  in  1810.  It  is  best 
prepared  by  adding  hydrochloric  acid  to  zinc  or  magnesium  tel- 

Zn  Te  +  2  H  Cl  =  H2  Te  +  Zn  Cla . 

It  is  a  colorless  gas,  which  entirely  resembles  the  hydrogen  com- 
pounds of  sulphur  and  selenium.  It  burns  readily,  with  a  blue 
flame,  and  is  gradually  decomposed  into  hydrogen  and  tellurium 
even  at  ordinary  temperatures.  It  is  instantly  changed  on  exposure 
to  the  air  :  -  H2  Te  +  0  =  H2  0  +  Te.t 

The  tellurides  are  with  difficulty  obtained  pure,  and  are  prepared 
like  the  sulphides  and  selenides. 

*  Brauner  ( Monatshefte  fur  Cheraie;  10,  411)  maintains  that  the  substance 
which  has  heretofore  been  regarded  as  tellurium  is  really  a  mixture ;  while  the 
pure  element,  if  subsequently  prepared,  will  probably  have  an  atomic  weight  of 
125  to  120,  yet  the  only  consistent  numbers  so  far  obtained  give  it  an  atomic 
weight  of  127.6. 

t  Compare  this  with  the  action  of  hydroiodic  acid  when  exposed  to  the 
air. 


ELEMENTS  OF  SULPHUR  FAMILY ;  TABLE  OF. 


105 


A  comparative  table  of  the  elements  of  the  oxygen  family  will 
serve  to  render  the  relationship  between  its  members  more  appar- 
ent, and  will  also  make  clear  the  resemblance  between  this  family 
and  the  halogens  :  — 

ELEMENTS. 


ATOMIC 
WEIGHT. 

SPECIFIC 
GRAVITY 
OF  SOLID. 

MELTING 
POINT. 

BOILING 
POINT. 

APPEARANCE. 

NOT- 
METALLIC 
PROP- 

ERTIES. 

o 

16. 





-182° 

Gas. 

s 

32.06 

2.04 

114° 

440° 

Yellow  solid. 

Se 

79.1 

4.5 

217° 

665° 

Dark    red   powder, 

black  when  fused. 

Te 

125.  (?) 

6.25  (?) 

500° 

above  1000° 

Silver  white,  metal- 

lic appearance. 

SPECIFIC 

SPECIFIC 

GRAVITY 

GRAVITY 

OF 

OF 

MOLECULE. 

VAPOR. 

VAPOR. 

AlR=l. 

H2  =  2. 

o 

1.1 

31.76 

02 

*  The  molecular  weights,  Se2  =  158,  Te2  = 

250,  are  somewhat  less  than  the  specific  gravi- 

ties found,  but  near  enough  to  show  that  the 

s 

2.2 

63.3t 

S2 

molecule  exists  as  two  atoms. 

t  The  specific  gravities  of  sulphur  and  se- 

lenium vapors  are  not  constant  below  1000°; 

they  gradually  become  larger  as  the  boiling 

Se 

5.7 

164.*t 

Se2 

points  are  approached.    The  molecules  S2  and 

i 

Se2  seem  to  form  larger  aggregations  as  the 

elements  approach  the  temperature  of  lique- 

faction ;  no  definite  formulae  seem  assignable 

Te 

9. 

259.* 

Te2 

to  these  molecules. 

HYDROGEN    COMPOUNDS. 


HEAT  OF 
FORMA- 

STABIL- 

HEAT  OF 
FORMA- 

STABIL- 

TION. 

TION. 

H,0 

684  K 

2  HF 

H2S 

H2Se 

27  K 
-  111K 

2HC1 
2HBr 

440  K 
242  K 

All  of  the  hydro- 
gen compounds  are 
colorless  gases  above 

H2Te 



2  HI 

-122K 

100°  Centigrade. 

106  ELEMENTS    OF   SULPHUR   FAMILY ;    TABLE   OF. 

On  comparing  the  atomic  weights  of  the  elements  of  the  oxygen 
family  with  those  of  the  halogens,  we  see  that  the  former  are, 
throughout,  somewhat  smaller  for  corresponding  elements ;  the  dif- 
ference, however,  is  but  slight* 

0      16  F     19 

S      32.06  Cl   35.45 

Se    79  Br  79.95 

Te  125  I     126.85 

*  Tellurium,  if  its  atomic  weight  is  127.6,  is  at  present  an  exception.     See 
page  104. 


VALENCE  AND  THE  OXYGEN  COMPOUNDS.       107 


CHAPTER   XVII. 

VALENCE  AND  THE  OXYGEN  COMPOUNDS  OP  THE  NOT-METALS. 

THE  elements,  the  properties  of  which,  we  have  studied,  form 
compounds  with  hydrogen,  all  of  which,  with  two  exceptions  — 
hydrogen  dioxide  and  the  corresponding  sulphur  compound  —  can 
be  obtained  as  vapors,  the  specific  gravities,  and  hence  the  molecular 
weights  of  which  can,  therefore,  readily  be  ascertained.  By  this 
means  we  arrived  at  the  conclusion  that  one  atom  of  chlorine,  bro- 
mine, or  iodine  could  unite  with  but  one  atom  of  hydrogen,  while 
one  of  oxygen  and  of  the  remaining  members  of  that  family  could 
unite  with  two.  These  elements,  therefore,  differ  among  themselves 
in  their  power  of  retaining  hydrogen  atoms.  In  addition  to  the 
foregoing  there  are  other  elements,  the  hydrogen  compounds  of 
which  are  composed  of  three  atoms  of  hydrogen  to  one  of  the  nega- 
tive element ;  these  elements  are  nitrogen,  phosphorus,  arsenic,  and 
antimony.  If  we  designate  any  one  of  these  elements  by  Y,  then 
the  formula  of  the  hydrogen  compounds  would  be  YH3 .  Only  two 
other  elements,  carbon  and  silicon,  form  gaseous  hydrogen  com- 
pounds; the  general  formula  of  these  is  ZH4  All  hydrogen 
compounds  can  therefore  be  classed  under  four  heads : 

YH,  XH2,  YH3,ZH4.* 

From  the  outset  we  have  considered  chemical  compounds  as  formed 
by  the  conjunction  of  the  atoms  of  elements  ;  the  atoms  themselves 
are  hypothetical, f  but  using  this  hypothesis  as  a  basis,  a  chemical 

*  We  can  imagine  all  apparent  variations  from  these  types  as  formed  by 
the  substitution  of  one  or  more  of  these  hydrogen  atoms  by  some  other  element, 
or  groups  of  elements.  Thus,  we  have  considered  sodium  hydroxide  as  water, 
in  which  one  atom  of  hydrogen  has  been  replaced  by  sodium ;  Na-O-H,  H-O-H ; 
hydrogen  dioxide  as  water,  in  which  one  atom  of  hydrogen  has  been  replaced 
by  hydroxyl  H-O— O-H,  H— O-H,  and  so  on. 

t  Sir  William  Thomson  considers  them  to  be  rings  formed  by  vortical 
motion  of  the  ether;  a  visible  example  of  such  motion  would  be  a  smoke  ring 
blown  by  a  locomotive. 


108  VALENCE;  HYDROGEN  COMPOUNDS. 

theory  productive  of  the  greatest  results  has  been  developed. 
"  Every  finite  quantity  of  matter  occupies  a  position  in  space  which 
is  definable  with  regard  to  other  material  particles;  the  question  as 
to  the  relative  position  (or  motion)  of  atoms  in  the  molecule  is 
scientifically  justified,  and  must  be  put  sooner  or  later  "  *  by  persons 
holding  the  atomic  theory.  This  problem  has  been  put  by  chemists 
ever  since  the  time  of  Berzelius ;  and  the  great  advance  in  organic 
chemistry,  which  has  been  reflected  in  inorganic  chemistry,  is  the 
result  of  its  successful  solution.  In  the  hydrogen  compounds  of  the 
not-metallic  elements  we  are  cognizant  of  the  number  of  atoms  in 
the  molecule,  because  we  have  been  able  to  determine  the  quantitative 
composition  of  these  compounds  and  also  the  molecular  weights. 

As  the  mass  of  the  atom  of  the  not-metal,  in  hydrogen  com- 
pounds, is  so  much  greater  than  that  of  the  hydrogen,  and  as,  in 
the  formation  of  new  chemical  compounds  from  those  of  hydrogen, 
it  is  always  the  hydrogen  which  is  replaced  by  some  other  element, 
provided  the  resulting  compound  remains  identical  or  similar  in 
character;  therefore,  it  is  more  than  probable  that  the  hydrogen 
atoms  are  joined  to  the  not-metal.  Whether  the  position  of  these 
is  fixed,  or  whether  they  are  free  to  rotate  around  the  not-metal,  we 
cannot  decide ;  but  recent  investigation  all  tends  toward  the  former 
theory.  In  the  compounds  VH,  XH2 ,  YH3 ,  and  ZH4 .  the  numerical 
capacity  possessed  by  the  not-metal  of  uniting  with  one,  two,  three, 
and  four  atoms  of  hydrogen  is  termed  the  valence  of  the  element, 
and  we  can  conveniently  express  this  valence  by  Roman  numerals 

I      II    III  IV 

placed  over  the  symbol  of  the  not-metal,  V,  X,  Y,  Z,  or  by  lines 
drawn  from  them,  as  V — ,  — X — ,  — Y —  or  — Z — .  The  element 

I 

which  can  unite  with  one  atom  of  hydrogen  is  said  to  be  univalent ; 
that  which  can  unite  with  two,  bivalent ;  with  three,  trivalent ;  and 
with  four,  quadrivalent;  hydrogen  is  always  univalent.  All  ele- 
ments, the  valence  of  which  is  more  than  one,  may  be  called  poly- 
valent. When  a  univalent  element  has  united  with  another  element 
or  radicle,t  its  capacity  for  further  union  has  ceased  ;  a  bivalent 
element,  however,  when  united  with  an  element  or  radicle  by  one 

*  W.  Ostwald,  Outlines  of  General  Chemistry,  Walker's  translation, 
t  A  group  of  elements  which  can,  as  a  whole,  replace  an  element  in  a 
chemical  compound,  is  called  a  radicle. 


VALENCE;  CHLORINE  COMPOUNDS.  109 

valence  has  not  lost  its  capacity  for  further  union  with  other  ele- 
ments or  radicles,  it  is  unsaturated;  for  example,  H — X —  is  in  this 

H 

I 
condition,  and  can  act  as  a  univalent  radicle.    Similarly  H —  Y —  is 

also  unsaturated  and  univalent;  H  —  Y — unsaturated  and  bivalent, 
H  H 

I  I  I 

H — Z — H,  H — Z — ,  and  H — Z — ,  unsaturated,  and  respectively  uni-, 

I          i  I 

bi-,  and  trivalent. 

Only  the  not-metals,  however,  form  hydrogen  compounds  obtain- 
able as  gases,  so  that  with  other  elements,  if  we  desire  a  similar 
means  of  determining  valence,  we  must  seek  for  gasifiable  compounds 
with  some  univalent  element  other  than  hydrogen.  Many  of  the 
metals  and  of  the  not-metals  are  capable  of  forming  such  compounds 
with  chlorine;  the  molecular  weights  of  these  can  therefore  be 
determined. 

The  halogens  form  compounds  with  formulse  analogous  to  those 
of  hydrogen,  and  such  compounds  can  obviously  be  used  to  deter- 
mine the  valence  of  the  elements ;  for  if  the  number  of  hydrogen 
atoms  with  which  an  element  is  united  in  a  molecule  indicates  the 
valence  of  that  element,  so  must  the  number  of  chlorine  atoms  in  a 
similar  molecule.  We  can,  therefore,  construct  a  table  containing 
a  series  of  chlorine  compounds,  just  as  we  did  with  the  hydrogen 
compounds,  and  further  investigation  shows  us  that  all  of  these 
compounds  can  be  brought  under  six  heads ;  using  M  as  a  general 
term  to  denote  an  atom  of  an  element  with  the  capacity  of  uniting 
with  chlorine,  we  have  :  — 

MCI,  MC12,  MCI.,  MC14,  MC16,  MC16,  and  using  Koman 
numerals  to  designate  the  valence :  — 

I          II       III       IV        V        VI 

M,    M,    M,    M,    M,    M. 

In  many  cases  in  which  the  hydrogen  compound  of  a  given  ele- 
ment exists,  we  can  also  study  the  chlorine  compounds,  so  that 
a  determination  of  the  valence  of  the  elements  by  means  of  the 
latter  offers  the  advantage  of  being  applicable  in  a  greater  number 
of  cases.  Sometimes,  as  is  the  case  with  some  metals,  the  bromide 


110       VALENCE  IN  COMPOUNDS  OF  POLYVALENT  ELEMENTS. 

or  iodide  is  obtainable  as  a  gas,  while  the  chloride  is  not ;  then  the 
former  compounds  answer  just  as  well  as  a  means  of  determining 
the  valence.  In  our  considerations  we  have,  so  far,  always  been 
able  to  appeal  to  experiment  to  answer  any  questions  which  may 
arise ;  in  those  which  follow  we  shall  have  to  indulge,  more  or  less, 
in  speculation. 

Can  elements,  polyvalent  toward  hydrogen,  unite  with  each 
other,  and,  if  so,  what  is  their  valence  ?  The  answer  to  the  first 
part  of  the  question  has  already  been  given ;  we  have  become  aware 
of  compounds  such  as  C02,  S02,  CS2,  all  of  which  are  formed  by  the 
union  of  polyvalent  elements ;  and  in  these  cases,  as  in  the  vast 
majority  of  those  which  fall  within  the  scope  of  this  book,  one 
atom  of  one  of  the  elements  in  the  molecule  unites  all  the  others ; 
furthermore,  substances  such  as  C02,  S02,  or  CS2  can  have  their 
molecular  weights  determined  in  the  same  way  as  can  those  of  com- 
pounds of  hydrogen,  so  that  the  same  reasoning  will  apply  with  the 
former  as  with  the  latter.  We  could,  therefore,  suppose  such  com- 
pounds to  have  a  structure  similar  to  that  of  water  (H — 0 — H)  ;  i.e., 
0 — C — 0,  0 — S — 0,  S — C — S  ;  in  such  an  event,  carbon  or  sulphur 
would  be  bivalent,  as  is  oxygen  in  water.  Indeed,  any  theory  other 
than  this  as  regards  the  valence  of  the  elements  in  those  compounds 
which  contain  bivalent  elements,  goes  beyond  the  realm  of  facts 
which  are  absolutely  proved  by  experiment.  In  spite  of  this,  the 
majority  of  chemists  have  thought  themselves  justified  in  holding 
other  views,  and  a  few  of  the  reasons  for  their  opinions  may  not  be 
out  of  place  here.  One  atom  of  oxygen  unites  with  two  atoms  of 
hydrogen  to  form  a  molecule  of  water,  so  that  in  this  compound  it 
is  undoubtedly  bivalent ;  furthermore,  oxygen  is  capable  of  uniting 
with  two  atoms  of  any  other  univalent  element,  such  as  sodium  or 
potassium,  the  oxides  of  which  are  Na2  0,  K2  0,  and  the  element  can 
also  unite  two  univalent  radicles,  or  groups  of  elements,  to  form 
compounds  (calling  any  radicle  q)  of  the  formula  q — 0 — q ;  so  that 
an  atom  of  this  element  can  serve  as  a  link  between  groups  of  ele- 
ments, a  function  which  is  evidently  impossible  for  univalent  ele- 
ments. When  oxygen  replaces  hydrogen  in  the  compounds  of  that 
element,  then  one  atom  of  the  former  always  takes  the  place  of  two 
of  the  latter ;  for  instance,  the  compound  CH4 ,  in  being  oxidized, 

forms,  in  addition  to  water,  C  j  Q2  at  first,  and  then  C  -j  Q  :  - 


VALENCE   OF   OXYGEN.  Ill 

CH4      +  20  =  CH20  +  H20;  arid 
CH2  0  H-  2  0  =  C02      +  H2  0. 

The  valence  of  an  element  can,  as  we  have  seen,  also  be  discovered 
by  a  study  of  the  formula  of  its  chloride ;  and  when  a  chloride  is 
converted  into  an  oxide,  one  atom  of  oxygen  always  replaces  two  of 
chlorine :  — 

CHLORIDES.  OXIDES.  CHLORIDES.  OXIDES. 

2NaCl  Na20  2A1C18  A1203 

2KC1  K20  2FeCls  Fe2O3 

CaCl2  CaO  C  C14  C  02 

FeCl2  FeO  2PC13  P2O3.* 

From  these  considerations  it  is  supposed  that  oxygen  remains 
bivalent  wherever  it  enters  into  chemical  combination.  In  assum- 
ing this  to  be  the  case,  we  must  consider  an  atom  of  oxygen  as 
united  to  other  elements  in  a  way  unlike  that  in  which  one  of  hy- 
drogen unites.  Applying  what  we  learned  in  regard  to  univalent 
elements,  where  we  saw  that  when  one  atom  of  an  univalent  ele- 
ment is  united  to  one  of  any  other,  its  further  power  of  union  is 
exhausted,  we  construct  the  following  arbitrary  rule :  — 

One  valence  of  any  element  in  a  chemical  compound  always  calls 
for  and  neutralizes  a  corresponding  valence  in  the  other  element  or 
elements  with  which  it  is  united. 

The  two  valences  of  a  bivalent  element,  therefore,  are  supposed 
to  neutralize  two  corresponding  valences  in  any  element  or  com- 
pound with  which  it  is  united.  The  following  examples  will  serve 
to  make  this  meaning  more  clear :  — 


C\H 
\H 

{Two  atoms  of  uni-  ^  C  Four  atoms  of  uni- 
valent hydrogen  re-  valent  hydrogen  re- 
placed by  one  of  bi-  1  I  placed  by  two  of  bi- 
valent oxygen,  carbon  f  |  valent  oxygen,  carbon 
remaining  quadriva-  remaining  quadriva- 
lent. J  1^  lent. 

*  All  of  the  chlorides  in  this  list  have  been  obtained  as  gases,  their  molec- 
ular weights  and  formulae  are  certain;  the  first  six  corresponding  oxides  have 
not,  but  having  once  determined  the  atomic  weights  of  the  elements,  the  for- 
mulae of  such  compounds  follow  from  their  composition  by  weight. 


112  FORMULA   OF    OXIDES. 

Where  an  atom  of  a  polyvalent  element  has  an  odd  number  of 
valences,  it  follows  that  these  cannot  be  exactly  neutralized  by 
those  of  an  atom  of  a  bivalent  element;  this  fact  will  become 
apparent  by  a  study  of  the  following  formulae  :  — 

n/Cl  ,0 

P  —  Cl   with   oxygen   yields    P          (phosphorus    remaining   triva- 

•\ci 


lent)  this    group  will   change   to   j       ^        with  one   oxygen   atom 

unsaturated  ;  and  the  latter  must  therefore  unite  with  some  other 
element  or  group  of  elements,  so  that  :  — 

° 


(phosphorus  remaining  trivalent  and  oxygen  uniting  the  two  univa- 
lent  groups  of  atoms).  By  a  similar  application  of  the  rule  we  can 
come  to  the  conclusion  that,  where  five  oxygen  atoms  unite  with  two 
of  some  other  polyvalent  element,  the  latter  is  quinquivalent,  and 
where  three  unite  with  one  it  is  hexavalent,  so  that  the  following 
table  of  the  formulae  of  oxides  can  be  constructed  (using  X  to  de- 
note an  atom  of  any  element)  :  — 

i 

X  20,  valence  of  X  one  ;  denoted  by  X2  0 

ii 


X  0,      « 

"  X  two  ; 

« 

«    X  0 

III 

X203,     « 

"  X  three  ; 

a 

"    X203 

IV 

X  02,     « 

"  X  four; 

tt 

"    X  02 

V 

X205,     « 

«  X  five  ; 

tt 

"    Xs05 

VI 

X  03,     « 

"  X  six; 

a 

"   X  03 

VII 

X207,     « 

"  X  seven  ; 

rt 

<e   X  0 

VIII 

X  04,     « 

"  X  eight  ; 

a 

«    X  04 

The  formulae  of  all  oxides,  with  a  few  exceptions,  correspond 
to  these  general  formulae,  and  what  has  been  said  of  oxygen  applies 
to  all  other  bivalent  elements.  In  endeavoring,  therefore,  to  ascer- 


FORMULAE    OF    OXIDES.  113 

tain  the  valence  of  an.  element  forming  an  oxide,  we  must  consider 
all  of  the  oxygen  atoms  in  the  molecule  to  be  held  in  position  by 
the  atoms  of  that  element,  just  as  is  the  case  in  the  structure  of  the 
hydrogen  and  chlorine  compounds  —  we  must  ascertain  the  gravi- 
metric composition  by  analysis,  so  that  we  can  learn  the  number  of 
atoms  united  in  the  formula  weight  ;  where  we  can  obtain  the  com- 
pound as  a  gas  we  must  learn  its  specific  gravity,  and  by  this  means 
its  molecular  weight.  For  instance,  there  are  two  oxides  of  phos- 
phorus, in  one  of  which  62  parts  by  weight  of  phosphorus  unite 
with  48  of  oxygen,  in  the  other  62  of  phosphorus  unite  with  80  of 
oxygen.  From  the  study  of  many  compounds  of  oxygen  and  phos- 
phorus, we  have  concluded  that  the  atomic  weights  of  these  two 
elements  are  16  and  31  respectively,  hence  the  composition  by 
weight  of  the  first  of-  the  two  compounds  leads  us  to  a  formula 
P2  O3,  that  of  the  second  to  a  formula  P2  05.  Now  we  can  construct 
two  formulae  by  writing  out  the  individual  atoms,  grouping  them 
together  so  that  connecting  lines  will  express  the  valences,  as 
follows  :  — 

o=r-o-p  =  o, 


Such  diagrams,  constructed  of  chemical  symbols  expressing  a 
theory  regarding  the  manner  in  which  atoms  are  grouped  in  a 
chemical  compound,  are  called  structural  formulas.  Structural 
formulae  may  be  more  or  less  complete,  in  order  to  express  the  prev- 
alent ideas  in  regard  to  the  relative  positions  of  the  atoms  in  a 
molecule.  The  two  given  above  are  not  intended  to  show  more  than 
that  the  oxygen  atoms  are  all  united  to  those  of  phosphorus  by 
means  of  two  valences  apiece,  and  that  the  two  phosphorus  atoms 
are  united  to  each  other  by  means  of  the  interposed  oxygen  atom  ; 
other  formulae,  however,  might  be  constructed  illustrating  a  theory 
regarding  the  relative  position  of  all  of  the  atoms,  as  has  been  done 
with  compounds  of  carbon,  in  which  field  recent  investigation  has 
brought  chemistry  to  a  point  where  the  relative  positions  in  space 
of  atoms  forming  a  molecule  can  be  studied.  It  will  fall  within 
the  scope  of  this  work  to  use  only  structural  formulae,  such  as 
those  which  are  given  above. 

A  trivalent  element  cafci  unite  with  bivalent  elements,  with  other 
trivalent  elements,  or  with  a  quadrivalent  element;  it  can  unite 


114  STRUCTURAL  FORMULAE. 

three  univalent  radicles,  and  so  on ;  its  valence  remaining  three,  one 
valence,  in  uniting,  neutralizes  another,  as  in  the  case  of  uni-  and 
bivalent  elements;  we  also  could  develop  similar  theories  with 
the  other  polyvalent  elements. 

Such  are,  briefly,  the  rules  which,  for  the  sake  of  uniformity, 
are  used  by  the  majority  of  chemists.  In  applying  such  methods, 
chemists  are  apt  to  allow  themselves  to  forget  how  far  experiment 
has  gone,  and  thus  to  confuse  the  theory  and  its  application  with 
the  phenomena  of  nature.  The  terms  "  valence  "  or  "  bond,"  used 
to  express  the  means  of  union  between  atoms,  may  become  to  them 
material  things,  the  lines  used  on  paper  to  express  the  supposed 
manner  of  union  of  atoms  may  simulate  real  linkings  between  exist- 
ing atoms  ;  so  that,  at  the  present  time,  the  science  of  chemistry  is 
in  danger  of  being  discredited  by  a  too  dogmatic  conception  of  the 
application  of  these  methods ;  and  thus  one  of  her  greatest  achieve- 
ments of  recent  times,  the  theory  of  valence,  may  ultimately  prove 
a  serious  obstacle  in  the  path  of  her  development.  In  comparing 
two  compounds,  such  as  H-O-H  and  0-C-O,  we  cannot,  without  in- 
dulging in  speculations,  assert  more  than  that,  in  one  case,  oxygen 
has  the  capacity  of  retaining  two  atoms  of  hydrogen  in  a  molecule ; 
and  that,  in  the  other,  carbon  presents  the  same  relationship  toward 
oxygen ;  the  supposition  that  oxygen  is  retained  in  the  molecule  C02 
by  a  force  differing  in  its  manifestations  from  the  one  retaining 
hydrogen  in  H2  0  is  purely  gratuitous,  and  based  upon  the  action  of 
oxygen  and  hydrogen  in  entirely  different  compounds.  In  studying 
substances,  the  molecular  weights  of  which  we  have  not,  as  yet,  been 
able  to  determine,  as  is  the  case  with  P2  03 ,  experiment  can  go  no 
farther  than  to  show  us  the  composition  of  such  compounds  by 
weight ;  of  the  structure  and  size  of  the  real  molecules  and  the  man- 
ner of  union  of  the  atoms  in  these  molecules,  we  have  no  knowledge, 
and  therefore,  with  such  bodies,  our  rules  of  valence  can  be  applied 
only  to  the  formula  weights.  If,  however,  we  remember  where  the 
boundary  between  speculation  and  experiment  lies,  we  shall  very 
much  simplify  the  study  of  chemistry  by  the  application  of  these 
rules.  The  valence  of  the  atom  of  an  element  is,  as  much  as  any 
other  property  of  that  element,  determined  by  the  family  in  which 
that  element  belongs,  and  hence  by  its  atomic  weight.  The  position 
of  an  element  in  the  periodic  system  oi.  the  elements,  therefore, 
offers  one  of  the  best  means  of  determining  its  valence.  An  atom 


ANHYDRIDES    AND    ACIDS. 


115 


may,  under  differing  conditions,  vary  the  number  of  oxygen  atoms 
with  which  it  unites,  thereby  changing  its  valence ;  the  not-metals 
are  especially  prone  each  to  form  a  number  of  oxides.  All  of 
the  not-metals,  with  the  exception  of  fluorine  and  bromine,  have 
oxides  which  are  constructed  according  to  certain  types.  The  fol- 
lowing is  a  table,  containing  the  formulae  of  those  oxides  which 
are  most  frequently  met  with ;  as  will  be  seen,  the  same  general 
structure  often  belongs  to  compounds  in  more  than  one  family.  X, 
in  each  family,  is  used  to  represent  a  typical  element :  — 


HALOGENS. 

OXYGEN  FAMILY. 

NITBOGEN  FAMILY. 

CASBON  FAMILY. 

I 

IV 

I 

IV 

X20 

X0.2 

x.2o* 

X02 

III 

VI 

III 

X203 

X03 

x.2o3 

v 

V 

X205 

X205 

VII 

X,07 

• 

Highest  valence 
seven. 

Highest  valence 
six. 

Highest  valence 
five. 

Highest  valence 
four. 

When  the  oxide  of  a  not-metal  is  the  anhydride  of  an  acid,  and 
is  changed  to  the  acid  by  the  addition  of  water,  the  valence  of  the 
atoms  of  the  element  which  characterizes  that  anhydride  is  not 
supposed  to  be  changed;  such  a  change  can  only  be  effected  by 
either  oxidation  or  reduction.  One  or  two  examples  will  make  this 
apparent.  C12  Of  is  a  compound  in  which  the  two  atoms  of  chlo- 
rine are  united  to  one  of  oxygen ;  chlorine  being  univalent  and  the 
structural  formula  of  this  oxide  comparable  with  that  of  water :  — 


H— 0— H 

Water. 


Cl— 0— Cl 

Chlorine  monoxide. 


When  chlorine  monoxide  is  added  to  water  the  following  change 
takes  place :   ^  +H20=2HOC1 

Chlorine  monoxide  -{-  Water  =  Hypochlorous  acid. 

The  nature  of  this  reaction  becomes  more  apparent  if  it  is  written 
as  follows :  — 

*  This  oxide  of  nitrogen  does  not  form  an  acid  on  addition  of  water,  but 
a  corresponding  acid  is  known,  which  breaks  down  into  water  and  this  oxide, 
t  X2  O  under  the  head  of  halogens  in  the  above  table. 


116  FORMATION   OF   CHLORINE  ACIDS. 

H— 0— H  H— 0     /    H 

H20  +  C120=          +  /    /   / 

Cl— 0— Cl  Cl     /   0— Cl 

Whether  we  consider  H — 0 — Cl  as  water  in  which  the'  atom  of 
hydrogen  is  replaced  by  one  of  chlorine,  or  as  chlorine  monoxide  in 
which  one  atom  of  chlorine  is  replaced  by  one  of  hydrogen,  is  a 
matter  of  indifference,  —  chlorine  obviously  remains  univalent  in 
the  acid,  as  it  was  in  the  anhydride.  Now,  in  order  to  change 
C120  to  C1203  we  must  add  oxygen.* 

Cl— 0— Cl  +  20     =     0  =  Cl— 0— Cl  =  0. 

The  process  of  oxidation  leaves  the  group  of  atoms  Cl — 0 — Cl 
intact.  Such  a  group  as  this  is  entirely  independent  of  any  added 
valence  assumed  by  chlorine.  Now,  when  C1203  reacts  with 
water :  -  ^  Og  -f  H2  0  =  2  H02  Cl 

Chlorine  trioxide  +  Water  =  Chlorous  acid,  or 

H— 0— H  H— 0 

H20  +  C1203  =  + 

0  =  Cl— 0— Cl  =  0          0  =  Cl       0— Cl  =  0  ; 

the  valence  of  chlorine  remains  three,  for  the  added  oxygen  atoms 
have  not  changed  the  nature  of  the  reaction.  What  is  true  of 
C120  and  C1203  must  be  true  of  X20  or  X203,  and  by  similar 
structural  formulae  we  can  show  that  when  the  reactions 

X205  +  H20  =  2  H03X,  and  X207  +  H20  =  2  H04X 

take  place,  the  valence  of  X  remains  unaltered,  for  the  point  of  at- 
tack for  the  water  is  the  oxygen  uniting  the  two  univalent  atoms, 
or  groups  of  atoms.  With  the  other  class  of  anhydrides  given  in  the 
table  named,  i.e.,  those  formed  by  not-metals  having  an  even  num- 
ber of  valences,  the  addition  of  water  has  no  different  result.  As 
there  is  but  one  atom  of  the  typical  element  present  in  these  oxides, 
the  entire  molecule  of  water  must  add  itself  to  one  molecule  of  the 
oxide ;  but  a  little  consideration  will  show  us  that  this  change  is 
identical  with  those  we  have  discussed,  for  an  oxygen  atom  which 
is  present  in  the  anhydride,  with  the  addition  of  a  molecule  of 
water,  forms  two  hydroxyl  groups :  — 

*  It  is  scarcely  necessary  to  remind  the  pupil  that  the  oxygen  atom  so 
added  must  be  fixed  by  chlorine;  the  oxygen  in  Cl — O — Cl,  being  bivalent, 
is  incapable  of  further  union. 


FORMATION   OF   SULPHUR   ACIDS. 


117 


a. 


a.    S02  +H20  =  H2S03 

Sulphur  dioxide  +  "Water  =  Sulphurous  acid. 

S03  +  H20=H2S04 

6.     Sulphur  trioxide  -f-  Water  =  Sulphuric  acid. 

H 

0—  H 
0—  H 


0 

II 
S 

II 
0 

0 

II 

b.   0=S 


+ 


H 


0 


H 

I 
0— H 


0—  H 
0  =  S—  0—  H 


0  0 

When  S02  is  changed  to  S03  by  oxidation,  the  valence  of  sulphur 
is  changed,  but  obviously  the  added  oxygen  atom  plays  no  more 
part  in  the  reaction  which  takes  place  by  addition  of  water,  than 
it  does  in  the  cases  previously  considered.  The  water  which  is 
taken  up  by  anhydrides  is  therefore  decomposed  by  them,  it  breaks 
down  into  H  and  -0-H.  The  hydrogen  attaches  itself  to  oxygen, 
while  the  hydroxyl  group  unites  with  the  atom  of  the  not-metal, 
which  characterizes  the  anhydride.  Water,  therefore,  does  not  enter 
into  the  composition  of  the  acids,  but  hydroxyl  groups  do.  In  the 
following  table  are  placed  the  general  formulae  of  the  oxides  pre- 
viously considered  and  the  corresponding  acids ;  X,  as  heretofore, 
denoting  the  not-metal :  — 

(     ' 

X20  +  H0  =  2lHO  X 


Ill 

X203 

v 

X205 

VII 

X207 

IV 

X  02 

VI 

X  0, 


_  9 


m 


H20  =  2VH03X 

(        VII\ 
+  H20  =  2^H04Xy 

IV 

+  H20  =  H203X 

VI 


oar 


118  FORMATION   OF   OXIDES   AND   ACIDS. 

The  formulae  of  these  six  oxides  and  of  the  corresponding  acids 
occur  more  frequently,  and  are  of  greater  importance,  than  any 
others  which  the  pupil  will  encounter  during  his  study  of  the  not- 
metals.  The  addition  of  oxides  of  metals  other  than  hydrogen 
to  the  anhydrides  must  necessarily  produce  reactions  similar  to 
those  we  have  just  considered. 


OXIDES   AND    OXY-ACIDS    OF    CHLOKINE.  119 


CHAPTER   XVIII. 

THE  COMPOUNDS  OF  CHLORINE  WITH  OXYGEN,  AND  WITH 
OXYGEN  AND  HYDROGEN. 

CHLORINE  forms  the  following  compounds  with,  oxygen,  and 
with  oxygen  and  hydrogen :  — 

1.  C12  0,  chlorine  monoxide,       1.     HO  Cl,  hypochlorous  acid, 

2.  C12  03 ,  chlorine  trioxide,        2.     H02  Cl,  chlorous  acid, 

3.  (Cl  02,  chlorine  dioxide,)*     3.     H03  Cl,  chloric  acid, 

4.     H04  Cl,  perchloric  acid. 

The  first  and  second  of  these  oxides  are  respectively  the  an- 
hydrides of  hypochlorous  acid  and  of  chlorous  acid ;  the  third  does 
not  form  any  corresponding  acid.*  The  anhydrides  of  the  third 
and  fourth  acids  are  not  known ;  whenever  an  attempt  is  made  to 
isolate  them,  complete  decomposition  takes  place.  The  formulae 
of  the  compounds  would  be  C12  05  and  C12  07 ,  were  they  capable  of 
existence. 

Where  there  are  four  acids  which  contain  oxygen,  and  which 
are  derived  from  the  same  element,  the  nomenclature  given  in  the 
following  table  has  been  adopted  ;  beginning  with  the  one  contain- 
ing the  least  oxygen,  we  have  :  — 

1.  Prefix  hypo  —  termination  in  —  ous. 

2.  No  prefix-  «  "  -ous. 

3.  No  prefix  —  "  «  -  ic. 

4.  Prefix  per-  "  "  —ic. 

Where  but  two  acids  are  known,  the  nomenclature  is  the  same 
as  under  2  and  3. 

Salts  derived  from  an  acid  terminating  in  -ous,  have  their 
names  end  in  -ite;  those  from  an  acid  terminating  in  -ic,  in 
—  ate.  The  prefix  in  the  name  of  the  acid  is  always  retained  in 
the  name  of  the  salt.  The  salts  obtained  from  the  above  acids  by 

*  On  being  passed  into  a  solution  of  potassium  hydroxide,  it  forms  potas- 
sium chlorite,  K  Cl  O2,  and  potassium  chlorate,  K  Cl  O3. 


120  CHLOEINE   MONOXIDE. 

replacing  the  hydrogen  with  some  other  metal  would  consequently 
be  named  as  follows,  with  potassium  as  the  metal  :  — 

1.  KO  Cl,  potassium  hypocbjorite, 

2.  K02  Cl,  potassium  chlorite, 

3.  K03  Cl,  potassium  chlorate, 

4.  K04  Cl,  potassium  perchlorate. 

Where  there  are  but  two  acids,  the  nomenclature  is  the  same  as 
under  2,  and  3. 

The  oxides  of  chlorine  all  are  unstable,  some  of  them  even 
extremely  explosive,  so  that  chlorine  and  oxygen  cannot  be  brought 
to  unite  directly.  In  endeavoring  to  find  a  cause  for  this  extremely 
ready  decomposition,  we  observe  that  these  oxides  are  produced  by 
the  union  of  two  very  similar  elements,  so  that  their  instability  is  a 
circumstance  which  is  certainly  not  unexpected  ;  for  oxygen  and  flu- 
orine, which  resemble  each  other  even  more  than  do  oxygen  and  chlo- 
rine, are  incapable  of  combination.  One  might  be  tempted,  from 
such  a  circumstance,  to  argue  that  not-metals  which  are  most  like 
each  other  would  have  the  least  chemical  affinity  ;  but  this  is  true 
only  in  a  very  few  cases.  The  oxides  of  sulphur  and  of  phosphorus, 
for  instance,  are  extremely  stable  bodies;  the  different  halogens  can, 
in  some  instances,  combine  with  each  other;  while  the  existence  of 
molecules  such  as  those  of  the  elements  (for  example  of  chlorine  and 
of  hydrogen),  shows  us  that  even  the  atoms  of  the  same  element 
may  be  most  firmly  united.  Any  attempts  to  generalize  regarding  the 
relative  stability  of  these  oxides  are  at  the  present  time  premature. 

(Chlorine  monoxide  is  not  of  itself  important,  but  the  acid  formed 
from  it  by  the  addition  of  water  is  of  the  greatest  commercial  value  ; 
its  salts,  more  especially  that  of  calcium,  being  the  compounds  most 
often  used  for  the  purpose  of  bleaching.  Chlorine  monoxide  can  be 
prepared  by  passing  chlorine  over  mercuric  oxide. 


It  is  a  reddish-yellow  gas,  with  a  penetrating  odor  resembling 
that  of  chlorine.  At  low  temperatures  it  is  condensed  to  a  red 
liquid,  which  boils  at  19°-20°  centigrade.  It  has  a  specific  grav- 
ity, air  being  1,  of  3.00,  hydrogen  being  2,  of  86.4.  Its  molecular 
weight  is  therefore  86.9,  its  formula  C12  0.*  It  gradually  dis- 

*  When  slightly  warmed,  chlorine  monoxide  explodes;  this  decomposition 
may  even  take  place  at  ordinary  temperatures. 


HYPOCHLOBITES.  121 

solves  in  water,  forming  an  orange-colored  solution  which  contains 
hypochlorous  acid :  - 

H  — 0  — H         H  — 0       H 

+  /  +  / 

Cl  — 0  — Cl        Cl      0  — Cl 

The  salts  derived  from  hypochlorous  acid  are  easily  prepared ; 
they  are  really  of  greater  importance  and  are  more  stable  than  is 
the  acid  itself. 

In  1788,  Berthollet  discovered  that  a  bleaching  solution  could  be 
prepared  by  passing  chlorine  into  a  solution  of  an  alkali.  He  at 
once  attempted  to  utilize  his  discovery  commercially,  and  so  estab- 
lished a  factory  at  Javelles,  where  bleaching  water  was  prepared  by 
treating  a  solution  of  potashes  (potassium  carbonate)  with  chlorine  ; 
this  solution  then  contained  hypochlorous  acid,  and  was  called  eau 
de  Javelles.  At  a  later  date,  lime  took  the  place  of  potashes  and 
the  resulting  bleaching  powder  became  known  as  chloride  of  lime. 

If  chlorine  is  passed  into  a  solution  of  potassium  hydroxide,  the 
first  product  will  necessarily,  because  of  the  great  chemical  af- 
finity between  the  metal  and  the  not-metal,  be  potassium  chloride ; 
therefore  :  --  K_  0  _H  +  Cl  =  K  Cl  +  —  0—  H. 

VV    G     \\    t"  Q*r  ^ 

The  group  — 'O  —  H   is,  however,  unsaturated,   so  that  it  can 

unite  with  another  atom,  in  this  case  with  one  of  chlorine,  H —  0  — 
-|-  Cl  =  H  —  0  —  Cl.  The  first  products  obtained  by  passing  chlo- 
rine into  a  solution  of  potassium  hydroxide  would  therefore  be  potas- 
sium chloride  and  hypochlorous  acid.  Now,  potassium  hydroxide 
is  a  base  ;  therefore,  with  hypochlorous  acid,  it  will  form  a  salt 
(potassium  hypochlorite),  and  water,  as  follows :  — 

KOH  +  HO  Cl  =  KO  Cl  +  H2  0. 

,  As  a  consequence,  the  ultimate  products  of  the  action  of  chlorine 
upon  potassium  hydroxide  are  potassium  chloride,  potassium  hypo- 
chlorite, and  water.  The  entire  change  is  summed  up  in  the  follow- 
ing reaction :  — 

2  KOH  4-  2  Cl  =  K  Cl  +  KO  Cl  +  H2  0.* 

*  That  free  hypochlorous  acid  is  at  first  produced,  and  that  this  acid  sub- 
sequently reacts  with  potassium  hydroxide  to  form  potassium  hypochlorite,  is 
proved  by  the  fact  that  if  a  weak  base,  which  is  incapable  of  forming  a  hypo- 
chlorite (zinc  oxide,  mercuric  oxide),  is  used,  then  the  chloride  and  hypochlo- 
rous acid  are  the  products  of  the  reaction.  See  also  Remsen,  Chemistry,  p.  113. 


122  HYPOCHLOROTJS  ACID;   DECOMPOSITION. 

If  calcium  hydroxide  is  used  in  place  of  potassium  hydroxide, 
then  calcium  chloride,  calcium  hypochlorite,  and  water  are  produced  ; 
the  formulation  of  the  reaction  would  be  modified  by  the  bivalence 
of  the  atoms  of  calcium,  which  have  twice  the  capacity  for  repla- 
cing hydrogen  in  acids  that  those  of  potassium  have.  This  reac- 
tion is  therefore  as  follows  :  — 

—  0—  H  (  —  01      H  —  0  —  Cl 

+  4Cl  =  Ca-3  +  and 

-0  —  H  (  —  01      H  —  0  —  01 

_  0—  H     HOC1  (_0  —  C1      HOH 

2.          Ca^  +  =Ca-}  + 

—  0  —  H     HOC1  (  —  0  —  01      HOH 

uniting  1  and  2,  we  have  :  — 

2  Ca  (OH),  +  4  01  =  Ca  C12  +  Ca  (0  Cl)2  +  2  H2  0.* 

A  moderate  heat  changes  hypochlorites  into  chlorates  ;  and  there- 
fore the  above  reactions  must  be  performed  in  the  cold  ;  the  one 
using  potassium  hydroxide  takes  place  only  with  dilute  solutions. 
The  hypochlorites  bleach  because,  when  acidified,  they  liberate 
chlorine.  The  first  change,  upon  addition  of  acids,  is  the  produc- 
tion of  hypochlorous  acid,  as  follows  t  :  — 

a.  K001     +    HOI      =KC1      +HOC1, 

b.  2  KOC1     -f    H2  S04  ==  K2  S04  +  2  01  OH, 

c.  Ca(OCl)2  +  2HCl      =CaCL>    +  2HOC1. 

This  portion  of  the  reaction  is  like  the  formation  of  hydrochloric 
acid  from  the  chlorides.  However,  hypochlorous  acid  is  unstable, 
and  breaks  down  as  follows  :  — 


HOC1  = 

and  hydrochloric  acid,  in   the  presence  of  nascent  oxygen,  forms 
*  There  is  some  reason  to  suppose  that  chloride  of  lime  contains  the  com- 

/   _  /-n 

pound   Ca  j  _  Q  ^j   a  substance  which  would  be  partly  chloride  and  partly 

hypochlorite;   the   relative   proportions   of  chloride   and  hypochlorite  would 

Cl 
remain  the  same  as  that  given  above,  for  2  Ca  _  Q  C1  would  contain  the  same 

percentage  of  calcium,  chlorine,  and  oxygen  as  Ca  C1.2  +  Ca  (OC1)2. 

t  A  solution  of  hypochlorous  acid  in  water  may,  if  dilute,  be  kept  for 
some  time. 


HYPOCHLOKOUS   ACID;   DECOMPOSITION.  123 

chlorine  and  water  (see  pages  59,  74)  ;  so  that,  if  hydrochloric  acid 
has  been  used  to  liberate  hypochlorous  acid,  chlorine,  and  not 
oxygen,  will  be  set  free  :  — 

H— N     Cl          H  Cl 

d.  HOC1  +  HC1  =  H20  +  2C1;  +   \          =      \     +     | 

H— 0— \  Cl     H— 0       Cl 

Now,  a  chloride  is  always  formed  simultaneously  with  a  hypo- 
chlorite  if  chlorine  acts  upon  such  bases  as  potassium  or  calcium 
hydroxide ;  so  that  where  sulphuric  acid  is  employed  to  liberate 
hypochlorous  acid,  the  following  changes  take  place :  — 

KOC1  +  KC1  +  H2S04  =  K2S04  +  HOI  +  HOC1; 
Ca(OCl)j  +  CaCl2  +  2H2S04  =  2CaS04  +  2HC1  +  2HOC1; 

after  they  are  set  free,  the  hypochlorous  acid  and  hydrochloric  acid, 
can  react  as  in  equation  d-,  the  complete  change  would  therefore 
be  represented  by  the  equations :  —  ^fc 

KOC1  +  KC1  +  H2S04  =  K2S04  +  H20  +201; 
Ca(OCl)2  +  CaCl2  +  2  H2S04  =  2  CaS04  +  2  H20  +  4  Cl ; 

Chlorine,  consequently,  is  liberated  when  hydrochloric  or  sulphuric 
acid  is  added  to  a  hypochlorite ;  and  the  acidified  hypochlorites 
exercise  their  bleaching  action  by  reason  of  the  liberation  of  that 
element.* 

Hypochlorous  acid  is  unknown  in  a  pure  state,  but  its  solution 
can  be  prepared  by  passing  the  anhydride  (C120)  into  water,  or 
better  still,  by  suspending  mercuric  oxide  in  water  and  then  sub- 
jecting it  to  the  action  of  chlorine ;  by  this  means  the  formation 
of  the  anhydride  is  avoided.  .  Concentrated  solutions  of  hypochlo- 
rous acid  possess  the  odor  of  chlorine  and  disintegrate  in  the  dark, 
although  they  change  more  rapidly  in  the  daylight.  If  the  acid  is 
quite  dilute  it  is  much  more  stable,  and  can  then  even  be  distilled 
without  great  decomposition.  A  solution  of  hypochlorous  acid  is 
an  energetic  oxidizer;  we  have  seen  that  hydrochloric  acid  is 
changed  to  chlorine  by  it,  and  other  hydrogen  compounds  are 

*  Calcium  hypochlorite  (bleaching  powder)  continually  gives  off  chlorine, 
when  it  is  exposed  to  the  air.  This  is  probably  due  to  the  action  of  the  car- 
bonic acid  of  the  atmosphere. 


124  POTASSIUM  CHLORATE. 

similarly  affected.  Hydrogen  sulphide,  for  instance,  is  acted  upon 
as  follows  by  hypochlorous  acid  :  — 

H2S  +  HOC1  =  S  +  H20  +  HCL 

Vegetable  dyes,  such  as  litmus  and  indigo,  are  instantly  bleached 
even  by  dilute  solutions  of  the  acid,  and  many  other  organic  sub- 
stances are  destroyed  by  it,  so  that  it  is  much  used  for  bleaching 
cotton  and  linen  goods.  With  silks  and  woollens  it  is  useless,  for 
these  it  colors  yellow. 

The  most  stable  acids  of  hydrogen,  oxygen,  and  chlorine  are 
those  which  contain  relatively  the  most  oxygen,  and  the  same  is 
true  of  the  corresponding  salts,  so  that  hypochlorous  acid  and  the 
hypochlorites  will  tend  to  change  into  compounds  with  more  oxygen 
in  proportion  to  the  same  amount  of  chlorine ;  one  portion  of  the 
salt  or  acid  being  oxidized  at  the  expense  of  the  other.  As  a  conse- 
quence, a  potassium  hypochlorite  solution  forms  potassium  chlorate 
when  it.  is  heated  to  boiling :  — 

3KC10  =  KC108  +  2KC1,  or, 
KC10 

KC10  KC1  0 

KC10  KCl^3' 

Now,  2  KOH  +  2  Cl  =  KC1  +  KC10  +  H20,  therefore, 

1.  6  KOH  +  6C1=3KC10  +  3KC1+3H20; 
and  when  the  solution  is  hot  and  concentrated, 

2.  3KC10  =  2KC1  +  KC103; 

therefore,  when  chlorine  is  passed  into  potassium  hydroxide  solution 
under  those  conditions,  the  result  (combining  equations  1  and  2),  is 
as  follows :  — 

6  KOH  +  6  Cl  =  5  K  Cl  +  K  Cl  03  +  3  H20. 

In  a  similar  way,  the  solution  of  calcium  hypochlorite  changes  to 
the  chlorate  of  calcium  on  heating.  Potassium  chlorate  is  much 
less  soluble  than  is  calcium  chlorate,  so  that  potassium  chlorate  can 
also  be  prepared  by  adding  the  solution  of  a  potassium  salt  to  a 


POTASSIUM   PERCHLORATE.  125 

solution  containing  calcium  chlorate ;   this  method  is  the  one  used 
for  preparing  the  salt  on  a  large  scale :  — 

Ca(C103)2  +  2KC1  =  2KC108  +  CaCl2. 

The  chlorates,  when  heated  to  a  temperature  considerably  higher 
than  that  required  to  effect  the  change  from  hypochlorites  to  chlo- 
rates, yield  oxygen  and  are  transformed  into  a  mixture  of  chloride 
and  perchlorate,  one  portion  of  the  salt  being  oxidized  at  the  ex- 
pense of  the  other :  — 


C1K 


(° 

KCUO- 
(0 

(° 

|-KCUO  = 
(0 

(oio 

i  KCllOjO 

(o;0 

KC10,  +  KC103  =  KC104  +  KC1  +  20.* 

Finally,  as  potassium  perchlorate  is  not  able  to  take  up  more  oxy- 
gen, it  breaks  down  into  potassium  chloride  and  oxygen  at  a  low 
red  heat  :  — 

KC104  = 


When  potassium  chlorate  is  heated,  the  salt  melts  at  a  moderate 
temperature  ;  when  the  latter  is  increased,  oxygen  begins  to  pass 
off;  the  salt  again  solidifies  when  it  has  changed  completely  to  a 
mixture  of  the  chloride  and  perchlorate  ;  finally,  it  once  more  melts 
at  a  red  heat  and  then  the  potassium  perchlorate  parts  with  all 
of  its  oxygen,  leaving  potassium  chloride  in  the  flask.  (See  page  20 
and  foot-note.) 

(The  chlorates,  especially  that  of  potassium,  are  but  little  infe- 
rior in  commercial  importance  to  the  hypochlorites.  They  are  used 
chiefly  for  their  oxidizing  powers,  while  potassium  chlorate  is  also 
of  medicinal  value.  In  using  a  chlorate,  care  must  be  taken  not  to 
have  the  salt  mixed  with  any  substance  easily  oxidized  ;  very  seri- 
ous explosions  have,  for  instance,  resulted  from  grinding  potassium 
chlorate  and  sugar  simultaneously  in  the  same  mortar.  Specimens 
of  chlorate  to  be  used  for  preparing  oxygen  should  always  be  first 

*  More  complicated  equations  are  frequently  given  for  this  reaction,  but 
although  there  is  some  variation  in  the  amount  of  oxygen  formed,  and  in  the 
relative  proportions  of  potassium  chloride  ahd  of  perchlorate  left  in  the  flask, 
yet  the  probability  seems  to  be  that  in  the  great  majority  of  cases  the  simplest 
equation,  which  is  that  given  above,  is  realized. 


126  CHLORIC   AND   PERCHLORIC   ACID. 

tested  on  a  small  scale,  in  order  to  insure  their  safety.  A  demon- 
stration of  these  facts  is  readily  supplied  by  rubbing  a  trace  of 
potassium  chlorate  over  a  very  small  bit  of  sulphur  in  a  rough 
mortar,  or  by  mingling  some  powdered  chlorate  with  red  phospho- 
rus, by  gently  brushing  the  two  substances  together  with  a  feather  ; 
when  the  mixture  is  struck  with  a  glass  rod  a  sharp  explosion  will 
result.*  The  commercial  application  of  potassium  chlorate  lies 
chiefly  in  the  preparation  of  fireworks  and  of  explosive  matches. 

Potassium  perchlorate  is  very  nearly  insoluble  in  cold  water, 
and  therefore  is  of  value  in  qualitative  analysis.  Because  of  its 
greater  stability  it  is  sometimes  used  for  pyrotechnic  purposes  in 
place  of  potassium  chlorate. 

Chloric  and  perchloric  acids  are  more  easily  decomposed  by  heat 
than  are  their  salts,  but  they  possess  greater  stability  than  do 
hypochlorous  or  chlorous  acids.  Either  can  be  prepared  according 
to  the  usual  method,  by  the  addition  of  sulphuric  acid  to  the  corre- 
sponding salt  :  - 

2KC103  +  H2S04  =  K2S04  +  2  HC103  1 


but  they  also  are  the  products  of  decomposition  of  those  chlorine 
acids  which  contain  less  oxygen;  these,  when  heated,  change  to 
chloric  acid,  but  as  hydrochloric  acid  is  formed  during  this  change, 
a  certain  amount  of  chlorine  must  also  be  given  off,  because  these 
powerful  oxidizers  always  destroy  such  compounds  of  hydrogen. 
Chloric  acid  finally  changes  to  perchloric  acid  upon  being  heated 
above  40°:  — 

2  H  Cl  03  =  H  Cl  04  +  H  Cl  +  2  0. 

Of  course,  the  oxygen  formed  by  this  decomposition  further  acts 
on  the  hydrochloric  acid,  forming  water  and  liberating  chlorine,  so 
that  the  reaction  is  more  complicated  than  the  equation.  (Of  all 
these  acids,  perchloric  acid  is  the  most  stable  ;  its  aqueous  solution 
can  be  distilled  without  decomposition,  so  that  by  this  means  it  can 
be  separated  from  the  sulphuric  acid  used  in  its  preparation. 
It  is  an  oily  substance,  and,  because  it  can  be  obtained  in  a  pure 

*  Only  very  small  quantities  must  be  used. 

t  It  is  better  to  use  barium  chlorate,  for  then  sulphuric  acid  will  form 
insoluble  barium  sulphate,  which  can  be  filtered  off. 


CHLORINE   TRIOXIDE;    CHLORINE   DIOXIDE.  127 

state,  best  illustrates  the  intense  capacity  for  oxidation  possessed 
by  these  chlorine  compounds.  This  is  shown  by  placing  a  drop  of 
the  acid  upon  paper  or  wood,  for  then  a  violent  explosion  ensues. 
The  acid  itself,  when  kept  for  some  time,  decomposes  spontaneously 
with  explosive  violence. 

Chlorine  trioxide,  chlorous  acid,  and  chlorine  dioxide  are  the 
only  chlorine  and  oxygen  compounds  which  remain  for  discussion. 

Chlorine  trioxide,  the  anhydride  of  chlorous  acid,  is  made  by  the 
reduction  of  chloric  acid  by  means  of  arsenic  trioxide.*  The  sub- 
stance is  a  green  gas  with  a  most  penetrating  and  irritating  odor. 
In  preparing  the  gas  the  temperature  must  be  kept  quite  low, 
otherwise  a  most  dangerous  explosion  may  result.  It  forms  a  dark- 
brown  liquid  at  the  temperature  of  snow  and  salt,  and  this  liquid 
decomposes,  even  if  it  is  kept  in  the  dark.  When  dissolved  in 
water,  chlorine  trioxide  produces  chlorous  acid  :  — 

C1203  +  H20  =  2HC102. 

The  solution  is  a  powerful  bleaching  and  oxidizing  agent.  On 
standing,  it  changes  to  chloric  acid  and  hydrochloric  acid,  which 
latter  is  further  oxidized  to  chlorine  and  water.  The  acid  neutral- 
izes bases  very  slowly.  The  salts  formed  by  such  neutralization  are 
called  chlorites,  and  are  powerful  oxidizers ;  many  of  them  change 
to  the  chlorates  quite  readily,  the  nature  of  this  alteration  being, 
in  principle,  the  same  as  that  accompanying  the  transformation  of 
chlorates  into  perchlorates.  Potassium  chlorite  is  converted  into 
the  chlorate  at  160°. 

Chlorine  dioxide  is  an  unstable,  greenish-yellow  gas,  formed 
when  concentrated  sulphuric  acid  acts  upon  potassium  chlorate. 
In  this  reaction  we  might  expect  the  production  of  chloric 
anhydride  :  -  2  H  C1  ^_^  Q  =  ^  ^  . 

for  the  concentrated  acid  would  remove  water  from  chloric  acid. 
This  is  not  the  case,  however,  for  the  chlorine  pentoxide,  which 
might  result,  is  incapable  of  existence,  so  that  a  part  of  its  oxygen 
is  used  in  oxidizing  chloric  acid  to  perchloric  acid  while  chlorine 
dioxide  is  given  off :  — 

HC103  +  C1205  =  HC104  +  2  C102. 

*  2  H  Cl  03  +  As2  03  =  H2  O  +  C12  O3  +  As2  O5.  The  H  Cl  O3  can  be  formed 
by  adding  nitric  acid  to  potassium  chlorate  :  K  Cl  O3  +  H  NO3  =  K  NO3  + 
H  Cl  03 ;  when  potassium  nitrate  and  chloric  acid  result.  The  existence  of 

p.hlorinfi  trirnrirlfi  is  donhtfnl.  '    SP.P.  T.iphio-'s  Anrmlpn    900  •   RJ. 


128  OXY-ACIDS   OF   CHLORINE ;    STRUCTURE. 

The  gas  is  a  most  powerful  oxidizing  agent ;  combustible  substances 
burn  in  it  with  explosive  violence.  This  may  be  shown  by  mixing 
some  potassium  chlorate  with  sugar  *  and  then  adding  a  drop  of  con- 
centrated sulphuric  acid,  when  the  mass  will  instantly  take  fire.  If 
a  little  chlorate  of  potassium  is  placed  in  a  deep  glass,  covered  with 
water,  a  small  piece  of  phosphorus  dropped  in  and  then  sulphuric 
acid  carefully  poured  directly  on  the  salt  by  means  of  a  pipette,  the 
combustion  of  phosphorus  by  means  of  the  chlorine  dioxide  liber- 
ated can  be  seen  to  take  place  under  the  surface  of  the  water.  If 
the  gas  is  warmed,  a  dangerous  explosion  results,  so  that  care  must 
be  taken  never  to  heat  a  mixture  of  sulphuric  acid  and  potassium 
chlorate.  The  specific  gravity  of  the  gas  shows  that  it  has  the  for- 
mula C102,  so  that  if  we  consider  oxygen  as  bivalent,  chlorine  is 
quadrivalent  in  this  compound. 

The  generally  accepted  theory  regarding  the  constitution  of 
these  acids  is  as  follows.  The  hydrogen  atom  is  not  attached  to 
chlorine,  but  to  oxygen,  forming  a  part  of  the  hydroxyl  group; 
and  this  hydrogen  atom  is  replaced  by  metals  when  the  salts  are 
formed  :  — 

H— 0— Cl  H— 0— Cl 

Hypochlorous  acid. 

H— 0— Cl  =  0 

Chlorous  acid. 

The  existence  of  the  hydroxyl  group  in  acids  containing  oxygen 
has  already  been  discussed.  (See  pages  115,  116,  117.) 

The  most  striking  characteristics  of  the  acids  composed  of 
chlorine,  oxygen,  and  hydrogen  are  —  their  intense  power  of  oxidiz- 
ing, their  extreme  instability,  and  the  tendency  which  those  with  a 
lesser  amount  of  oxygen  have  to  change  into  those  with  a  greater. 
Although  the  salts  are,  as  a  rule,  less  easily  decomposed,  they 
nevertheless  display  similar  properties. 

*  Do  not  rub  in  a  mortar. 


HYPOBROMOUS   ACID.  129 


CHAPTER   XIX. 

COMPOUNDS  OF  BROMINE  AND  OF  IODINE  WITH  OXYGEN  AND 
HYDROGEN,  THE  COMPOUND  OF  IODINE  WITH  OXYGEN, 
AND  THE  COMPOUNDS  OF  THE  HALOGENS  WITH  EACH 
OTHER. 

ALL  attempts  to  isolate  oxides  of  bromine  have  proven  futile; 
unstable  as  the  oxides  of  chlorine  are,  those  of  bromine  are  evi- 
dently still  more  so.  The  acids  containing  bromine,  oxygen,  and 
hydrogen  are  known  only  in  aqueous  solutions;  their  formulae 
correspond  to  those  of  the  chlorine-acids,  but  bromous  acid  is 
unknown,  and  the  existence  of  perbromic  acid  is  very  doubtful. 
The  only  compounds  of  bromine,  oxygen,  and  hydrogen  with  which 
we  have  to  deal  are  therefore  hypobromous  acid,  H  Br  0,  and  bromic 
acid,  H  Br  03 ,  while  in  addition  we  must  discuss  the  salts  derived 
from  these. 

Solutions  of  hypobromous  acid  in  water  are  produced  under  cir- 
cumstances exactly  analogous  to  those  which  were  observed  in  the 
preparation  of  hypochlorous  acid ;  such  solutions  have  powerful 
bleaching  properties,  and  are  very  readily  decomposed  even  by 
slight  warmth.  When  bromine  is  added  to  very  dilute  potassium 
hydroxide  solution,  a  liquid  having  bleaching  properties  is  pro- 
duced; the  reaction  is  similar  to  that  encountered  in  studying  the 
action  of  chlorine  on  a  dilute  and  cold  solution  of  caustic  potash. 
(See  page  121.) 

2  KOH  +  2  Br  =  KBr  +  KOBr  +  H20. 

When  the  potassium  hydroxide  solution  is  too  concentrated, 
bromate  of  potassium  is  produced  even  at  ordinary  temperatures, 
the  conversion  of  hypobromites  into  bromates  being  a  change  much 
more  readily  produced  than  the  corresponding  one  with  chlorine, 
but  the  principle  of  the  action  is  the  same :  — 

6  KOH  +  6  Br  =  5  K  Br  +  K  03  Br  +  3  H2O  * 

*  It  is  not  necessary  to  enter  into  the  explanation  of  the  course  of  these 
reactions;  the  pupil  should  undertake  this  by  repeating  the  various  phases 
given  under  chlorine,  while  substituting  bromine  for  the  latter  element. 


130  BKOMIC   ACID  ;   IODIC    ACID. 

The  bromate  of  potassium  is  not  very  soluble,  so  that  it  can  be  fil- 
tered from  the  solution  containing  the  bromide  and  then  be  recrys- 
tallized  from  hot  water ;  bromate  of  barium  can  be  prepared  in  a 
similar  manner.  The  latter  will  yield  a  solution  of  bromic  acid 
when  exactly  enough  sulphuric  acid  to  form  barium  sulphate  is 
added : — 

Ba  (Br  03)2  +  H2  S04  =  Ba  S04  +  2  H  Br  03 . 

The  solution  of  bromic  acid  is  colorless,  and  may  be  concentrated 
by  evaporating  the  excess  of  water  in  a  vacuum,  but  when  warmed 
the  acid  breaks  down  completely  into  bromine,  oxygen,  and  water. 
Naturally,  all  of  the  compounds  under  consideration  are  powerful 
oxidizers ;  the  bromates  form  very  explosive  mixtures  with  oxidiz- 
able  substances,  while  the  bromate  of  ammonium  may  even  explode 
spontaneously. 

Iodine  forms  the  pentoxide  I205,and  two  acids,  iodic  acid, 
HI08,  and  per-iodic  acid.  If  the  anhydride  of  the  latter  acid 
existed,  it  would  have  the  formula  I2  07 ,  for  in  the  previous  chap- 
ter we  saw  that  the  theoretical  anhydride  of  perchloric  acid  would 

VII 

be  C12  07 ;  now,  by  the  addition  of  water  to  these  anhydrides,  the 
first  products  would  be  per-iodic  and  perchloric  acids  respectively, 
as  follows  :  — 

I,07  +  H20  =  2HI04, 
C1207+H20=2HC104. 

If  we  recall  the  structural  formulae  of  these  acids,  it  seems 
reasonable  to  suppose  that  the  oxygen  atoms  contained  in  them 
would  be  capable  of  adding  the  elements  of  water  to  form  hydroxyl 
groups,  in  conformity  with  the  tendency  manifested  in  the  produc- 
tion of  such  groups  by  the  addition  of  water  to  the  anhydrides,  and 
in  this  way  more  complicated  compounds  would  result :  — 


=  0  +  H-OH 


=  0 


— U-H  ~  TT  — O-H 


— 0-H 


O-H 


—O-H 
—O-H 
—O-H 
—O-H 
—O-H 
—O-H 
—O-H 


First  acid,  Second  acid,  Third  acid,  Fourth  acid. 

X04H  +  H20  =  X  05  H3  +  H20  =  X  06  H6-  +H20  =  X07  H7 . 


PER-IODIC    ACID;    IODINE   PENTOXIDE.  131 

With  the  addition  of  each  molecule  of  water,  one  oxygen  atom 
of  the  acid  can  yield  two  hydroxyl  groups,  until  finally  all  have 
been  converted,  and  a  complete  hydroxide  has  been  produced. 
This  last  acid  is  called  the  normal  acid.  By  separating  water  from 
the  normal  acid,  the  various  other  acids  can  be  formed ;  so  that  in 
the  end  we  arrive  at  an  anhydride,  which  is  obviously  identical 
with  that  from  which  we  started.  None  of  these  changes  involve 
either  an  oxidation  or  a  reduction,  for,  if  such  were  to  take  place, 
we  should  produce  acids  derived  from  different  anhydrides,  in  which 
the  valence  of  the  characterizing  element  would  vary;  therefore,  we 
can  assume  that  in  the  change  from  the  anhydride  to  the  first  acid, 
and  in  the  subsequent  conversion  of  this  to  the  normal  acid,  no 
alteration  in  the  valence  of  X,  in  the  above  formulae,  has  taken 
place.  The  process  of  adding  water  to  these  anhydrides  is  called 
hydration,  and  the  acids,  excepting  the  ones  with  least  amount  of 
hydrogen,  are  called  hydrated  acids.  Hydrated  acids  occur  quite 
frequently,  but  normal  acids  are  extremely  unstable  ;  their  existence 
even  in  solution  is  doubtful,  for  when  a  large  number  of  hydroxyl 
groups  are  attached  to  the  same  element,  they  will  always  show  a 
great  tendency  to  separate  water ;  yet  often  the  acids  lying  between 
that  having  the  least  hydrogen  and  the  normal  acid  have  no  tend- 
ency to  break  down;  in  fact,  they  are  sometimes  the  only  ones 
which  we  encounter.  In  many  of  the  hydrated  acids  only  a  portion 
of  the  hydrogen  atoms  can  be  replaced  by  metals  to  form  salts,  a 
fact  not  surprising  if  we  consider  that,  as  soon  as  a  hydrogen  atom 
in  an  acid  is  replaced  by  a  more  metallic  element,  the  whole  com- 
pound is  rendered  more  positive,  and  therefore  has  its  tendency  to 
take  up  positive  elements  diminished.  Per-iodic  acid  exists  in  the 
hydrated  form  H5  06 1,  corresponding  to  the  third  acid  of  the  series 
given  on  the  table  above. 

The  pentoxide  of  iodine  is  a  white  powder,  which  melts  at  300°, 
and  then  instantly  decomposes  into  oxygen  and  iodine ;  it  is  pro- 
duced by  oxidizing  iodine  with  nitric  acid,  or  by  heating  iodic  acid 
for  some  time  at  170°.  The  oxides  of  chlorine  are  all  endothermic 
substances,  but  453  K  are  liberated  in  the  formation  of  I2O5,  so 
that,  while  the  former  compounds  are  explosive,  the  latter  is  quite 
stable.  From  this  we  see  that,  with  the  diminishing  not-metallic 
character  of  the  halogens,  there  appears  a  diminishing  stability  of 
the  compounds  of  those  elements  with  metals,  while  at  the  same 


132  IODIC  ACID;  PER-IODIC  ACID. 

time  an  increasing  stability  of  the  oxides  is  manifested;  but  when 
we  try  to  draw  general  conclusions  from  these  facts,  we  must 
remember  that  the  oxides  of  bromine  are  less  stable  than  those  of 
chlorine. 

Iodic  acid  can  be  prepared  by  oxidizing  iodine,  suspended  in 
water,  by  means  of  chlorine,  or  by  adding  the  anhydride,  I2  05 ,  to 
water.  It  is  a  crystalline  solid  with  powerful  oxidizing  properties  ; 
phosphorus  and  arsenic,  for  instance,  are  oxidized  by  it  respectively 
to  phosphoric  and  arsenic  acids,  and  it  even  changes  graphite  to 
carbon  dioxide. 

The  iodates  are  formed  either  by  adding  a  base  to  iodic  acid ;  — 

MOH  +  HI03 ,  =  MI03  +  HOH,* 

or  by  dissolving  iodine  in  an  alkali,  a  reaction  which  is  parallel  to 
that  which  takes  place  with  chlorine  or  bromine :  — 

6KOH  +  6I  =  5KI-f  KI08  4-  3  H20.      oj^ 

Iodic  acid  can  be  liberated  by  adding  a  non-oxidizable  acid  to 
the  iodates.  Many  of  the  iodates  form  per-iodates  when  heated, 
but  the  latter  can  also  be  produced  from  the  former  by  the  addition 
of  some  oxidizer.  Two  of  the  per-iodates  (AgI04,  KI04)  cor- 
respond to  the  perchlorates  in  formula,  but  the  majority  of  the 
salts  of  per-iodic  acid  are  derived  from  the  hydrated  acids ;  for  in- 
stance, Na5 106  from  H5 106 ,  and  Ag3  I05  from  H3 105 .  Salts  of 
per-iodic  acid,  in  which  only  a  part  of  the  hydrogen  atoms  have  been 
replaced  by  other  metals,  are  known ;  an  example  of  such  a  salt 
would  be  Na^  H3 106 .  Per-iodic  can  be  isolated  from  its  salts  by  the 
addition  of  some  other  acid,  and  when  separated  from  its  solutions 
by  slow  evaporation,  it  is  a  crystalline  solid  of  the  formula  H5 106 , 
and  is  a  powerful  oxidizer.  Other  more  complicated  per-iodic  acids 
and  per-iodates  exist,  but  for  their  study  the  pupil  must  be  referred 
to  some  larger  work. 

The  halogens  can  form  a  number  of  compounds  with  each  other. 
Three  of  these  are  produced  by  the  union  of  the  elements,  atom  for 
atom  ;  they  are  Br  Cl,  bromine  monochloride,  an  unstable  liquid 

*  It  obviously  makes  no  difference  whether  we  write  iodic  acid  HIO3  or 
HO3 1,  sulphuric  acid  H2  O4  S,  or  H2  SO4 ,  etc.,  excepting  in  cases  where  we  in- 
tend to  convey  some  idea  in  regard  to  the  structural  formulae  of  the  acids;  but 
the  method  of  writing  the  formulae  of  acids  with  the  symbol  of  oxygen  as  the 
terminal  letter  is  the  one  rendered  more  familiar  by  usage.  Both  systems 
are  employed  in  this  book. 


OXY-ACIDS    OF   THE   HALOGENS  ;    TABLE   OF. 


133 


decomposing  above  10° ;  I  Cl,  iodine  monochloride,  a  fluid  which  is 
readily  decomposed  by  water;  and  IBr,  iodine  monobromide,  a  more 
stable,  crystalline  solid.  In  addition  to  these,  a  solid  trichloride  of 
iodine,  I C13 ,  and  a  liquid,  pentafluoride  IF5 ,  are  known.  Com- 
pounds having  the  formulae  Br  Cl,  I  Cl,  and  I  Br  are  formed  by  the 
direct  union  of  the  elements  ;  IC13,  by  the  addition  of  chlorine  to 
I  Cl ;  and  IF5 ,  by  the  action  of  iodine  on  the  fluoride  of  silver.  All 
of  these  compounds  are  decomposed  by  the  addition  of  water, 
although  I C13  can  exist,  provided  but  little  water  is  present.  The 
reaction  with  I F5  is  as  follows  :  — 

IF5  +  3H20  =  HI03  +  5HF. 

The  decomposition  of  the  fluoride  and  the  chloride  of  iodine  by 
means  of  water  is  a  change  similar  to  those  produced  by  the  hydra- 
tion  of  other  halides  of  not-metals.  For  example,  phosphorus  tribro- 
mide  breaks  down  into  phosphorous  acid  and  hydrobromic  acid,  when 
it  is  added  to  wa'ter.  (See  page  80.)  lodic  acid,  which  is  derived 
from  an  oxide,  I2  05 ,  in  which  iodine  is  quinquivalent,  is  therefore 
also  formed  from  a  fluoride  of  iodine  in  which  the  element  is  like- 
wise quinquivalent.  In  comparing  the  formulae  of  I  Br,  I C13 ,  and 
I F5 ,  we  are  impressed  with  the  fact  that  the  more  atoms  of  an- 
other halogen  can  combine  with  an  atom  of  iodine,  the  greater  the 
difference  between  the  atomic  weights  of  the  two  uniting  elements. 
The  following  is  a  table  of  the  formulae  belonging  to  the  com- 
pounds discussed  in  the  last  two  chapters :  — 


CHLORINE. 

BBOMINE. 

IODINE. 

Oxides  C12  O 
C13  03 
Cl   O2 

Acids  HO  Cl 
HO2C1 

Acids  H  0  Br 

H03C1 
HO4C1 

H03Br 
H  04  Br* 

I205 

H03I 
H04IJ 

*  Existence  doubtful. 

t  The  hypoiodite  of  potassium  probably  exists  in  solution  immediately 
after  adding  iodine  to  a  cold  and  dilute  solution  of  potassium  hydroxide;  it, 
however,  soon  changes  to  the  iodate  on  standing.  Hypoiodous  acid  is  un- 
known. See  Colischonn ;  Zeitschrif t  f  iir  Analyt.  Cliem. ;  29,  566. 

J  This  acid  exists  in  its  hydrated  form,  H5IO6 .  The  stability  of  all  of  the 
acids  and  of  their  salts  increases  with  increasing  number  of  oxygen  atoms. 
They  are  all  powerful  oxidizers. 


134          OXIDES    OF   ELEMENTS    OF    THE   SULPHUR   FAMILY. 


CHAPTER   XX. 

THE   COMPOUNDS   OF    THE    ELEMENTS   OF    THE   SULPHUR  FAM- 
ILY  WITH   OXYGEN,  AND   WITH   OXYGEN   AND   HYDROGEN. 
SULPHUR  DIOXIDE  AND    SULPHUROUS  ACID. 

Sulphur  dioxide  and  sulphurous  acid.  Sulphur  dioxide — Formula, 
S02;  specific  gravity,  air  =  1,  is  2.23,  H2=2,  is  64.22  j  1  c.  c. 
at  0°  and  .76  m.  pressure  weighs  .002896  gram. 

THE  oxygen  compounds  of  the  elements  of  the  sulphur  family, 
while  they  resemble  those  of  the  halogens  to  a  certain  extent,  nev- 
ertheless differ  widely  from  the  latter,  both  in  their  formulae  and 
characteristics.  The  two  series  of  compounds  resemble  each  other 
chiefly  because  the  members  of  both  are  anhydrides  ;  they  differ  very 
greatly,  however,  in  the  ease  with  which  they  are  decomposed.  The 
oxides  of  chlorine  are  explosive  compounds,  while  that  of  iodine 
is  disintegrated  at  300° ;  on  the  other  hand,  the  oxides  of  the  ele- 
ments of  the  sulphur  family  are  quite  stable  and  have  different  for- 
mulae. The  oxides  of  the  halogens  (with  the  exception  of  Cl  02)  are 
formed  by  joining  two  atoms  of  the  not-metal  by  means  of  a  bi- 
valent oxygen  atom,  as  in  Cl  —  0  —  Cl  and  0  =  Cl  —  0  —  Cl  =  0, 
while  the  oxides  of  the  sulphur  group  have  no  such  linking,  as 
will  be  seen  from  the  formulae  of  the  following  compounds  :  —  * 

S  j  and    S    -5  =  0 

(  =  0 

Sulphur  dioxide.  Sulphur  trioxide. 

This  difference  in  the  formulae  of  the  anhydrides  produces  a  dif- 
ference in  those  of  the  acids  derived  from  them;  for,  in  changing 
the  oxides  of  the  halogens  into  acids,  the  water,  in  order  to  form 
two  hydroxyl  groups,  attacks  the  linking  oxygen,  and  thus  gives  us 
acids  containing  but  one  hydrogen  atom ;  but  in  converting  the 
oxides  of  the  sulphur  family  no  such  separation  can  take  place. 

*  Oxides  of  sulphur  with  the  formulae  S2O3  and  S2O7  have  been  de- 
scribed, but  these  are  comparatively  unimportant. 


OXY-ACIDS   OF   ELEMENTS   OF   THE   SULPHUR   FAMILY.      135 

The  following  formulae  will  make  this  more  apparent  (X  denotes 
a  halogen  atom  and  Y  an  atom  of  any  element  of  the  sulphur 
group)  :  - 

X  — 0       H  X  — OH  ,,(=0  /  =0 

and  1  X       ^   ,  TT          ^  \  Q H 


X      0  — H       X— OH  0  — H      (  — 0— H 

X2  0  +  H20  =  2X  OH  and  Y02  +  H2  0  =  Y03H2 

The  above  makes  it  plainly  evident  that,  while  the  first  acid 
derived  from  any  of  the  oxides  of  the  halogens  must  contain  one 
hydrogen  atom,  those  derived  from  the  sulphur  group  must  contain 
two.  As  a  general  rule,  the  most  important  acids  in  any  chemical 
family  have  as  many  replaceable  hydrogen  atoms  as  are  contained 
in  the  corresponding  hydrogen  compounds ;  for  instance,  H  Cl  and 
HC103,  HI  and  HI03,  H2S  and  H2S04,  H2Se  and  H2SeO3. 
All  of  the  hydrated  acids  must  bear  a  simple  relationship  to  these, 
for  they  are  formed  therefrom  merely  by  the  addition  of  water. 
(See  page  131.)  Subsequently  we  shall  see  that  the  same  rule  ap- 
pertains to  the  elements  of  the  nitrogen  family. 

The  elements  of  the  sulphur  family  form  the  following  oxides 
and  acids :  — 

502  sulphur  dioxide      +  H2  0  =  H2  S03     sulphurous  acid. 

503  sulphur  trioxide    -j-  H2  O  =  H2  S04     sulphuric  acid. 
Se  02  selenium  dioxide    -f  H2  0  =  H2  Se  03  selenious  acid. 
Te  02  tellurium  dioxide  +  H,  0  =  H2  Te  03  tellurous  acid. 
Te  03  tellurium  trioxide  +  H,  0  =  H2  Te  04  telluric  acid. 

Two  oxides,  S2  03  and  S2  07 ,  have  also  been  made,  while  a  num- 
ber of  sulphur  acids  of  less  importance  than  the  above  exist. 
Mention  of  these  will  be  made  at  the  proper  time.  All  the  ele- 
ments of  the  sulphur  group,  on  burning  in  air  or  oxygen,  form  the 
dioxides ;  and  these  can  be  converted  into  the  trioxides  by  oxida- 
tion, excepting  in  the  case  of  selenium,  the  trioxide  of  which  has 
never  been  prepared. 

The  natural  occurrence  of  sulphur  dioxide  is  limited  to  the  gases 
which  escape  from  the  craters  of  volcanoes ;  but,  as  sulphur  is  gen- 
erally found  in  the  coals  used  as  fuel,  sulphur  dioxide  must  be  a 
product  of  their  combustion,  and  hence  occurs  in  minute  traces  in 
the  atmosphere  of  cities,  although,  being  moist,  it  is  rapidly  ox- 


136          SULPHUR  DIOXIDE;  HISTORY,  PREPARATION. 

idized.  The  history  of  sulphur  dioxide  is  as  ancient  as  that  of 
sulphur  itself  ;  for,  as  it  is  produced  by  the  combustion  of  the  latter, 
its  properties  could  not  fail  to  become  an  object  of  interest.  The 
Romans  were  well  acquainted  with  the  disinfecting  powers  of  burn- 
ing sulphur  and  used  it  in  cleansing  their  wine-skins  ;  sulphur 
dioxide  was  confounded  with  sulphuric  acid  by  the  alchemists  ; 
Stahl  first  proved  its  individuality  ;  Priestley  obtained  it  pure  by 
collecting  it  over  mercury,  and  Lavoisier  explained  its  composition. 

The  preparation  of  sulphur  dioxide. 

Sulphur  dioxide  can  be  formed  either  by  oxidizing  sulphur  or  by 
deoxidizing  sulphuric  acid.  With  the  first  method  we  are  already 
acquainted,  for  we  saw  that  sulphur,  or  combustible  substances 
containing  sulphur,  yield  sulphur  dioxide  when  they  are  burned  ; 
the  second  method  is  best  employed  in  preparing  the  gas  for  labo- 
ratory use.  Substances  such  as  charcoal,  sulphur,  and  some  of  the 
metals,  will  reduce  sulphuric  acid  while  they  themselves  become 
oxidized.  If  charcoal  is  heated  with  sulphuric  acid,  carbon  dioxide 
and  sulphurous  acid  are  produced  :  — 

2  H2  S04  +  C  =  C02  +  2  H2  S08; 

but  the  latter,  being  an  acid  the  anhydride  of  which  is  a  gas,  breaks 
down  into  water  and  that  anhydride  :  — 


Sulphur  acts  in  a  manner  similar  to  charcoal,  with  the  difference 
that  with  it  only  sulphur  dioxide  can  be  formed  :  — 

2  H2  S04  -f  S  =  S02  +  2  H2  S03,  and  2  H2S03  =  2  H20  +  2  S02, 
so  that  :  — 

2  H2  S04  +  S  =  3  S02  +  2  H2  0. 

Better  than  either  of  these  methods  is  the  preparation  by  means 
of  copper  and  sulphuric  acid.  Cold  sulphuric  acid  has  very  little 
action  on  copper  ;  but  if  copper  shavings  are  heated  with  sulphuric 
acid,  sulphur  dioxide  will  be  given  off.4'2  Some  doubt  exists  as  to 
the  mechanism  of  this  reaction.  One  explanation  which  has  been 
offered  is  as  follows.  When  dilute  sulphuric  acid  acts  on  zinc, 
hydrogen  is  produced  :  — 

Zn  +  H2  S04  =  Zn  S04  +  2  H, 

but  if  the  sulphuric  acid  is  hot  and  concentrated,  not  hydrogen,  but 
sulphur  dioxide  is  formed.     It  is  therefore  reasonable  to  suppose 


SULPHUR  DIOXIDE;  PREPARATION.  137 

that  the  first  result  of  the  contact  of  zinc  and  sulphuric  acid  is 
always  the  liberation  of  hydrogen  ;  but  when  the  acid  is  hot  and 
concentrated  the  conditions  are  so  altered  that  it  will  give  up  its 
oxygen  very  readily.  The  hydrogen  which  is  being  generated  would 
then  form  Water  with  the  oxygen,  so  that  sulphurous  acid  would  be 
set  free :  — 

1.  Zn-fH2S04  =  ZnS04  +  2H 

2.  2  H  +  H2  S04  =  2  H2  0  +  S02 ,  combining  1  and  2  we  have, 

3.  Zn  +  2  H2  S04  =  Zn  S04  +  2  H2  0  +  S02 . 

Now,  although  copper  produces  no  hydrogen  with  sulphuric 
acid,  yet  it  can  be  conjectured  that  the  reaction  takes  place  in  a 
manner  similar  to  that  given  above,  substituting  copper  for  zinc. 
There  is  strong  reason  to  suppose,  however,  that  when  metals  pro- 
duce sulphur  dioxide  from  sulphuric  acid,  they  act  exactly  as  do 
carbon  or  sulphur,  by  removing  oxygen  without  having  to  call  in 
the  aid  of  hydrogen  :  — 

H,  S04  +  Cu  =  H2  S03  +  Cu  0. 

The  copper  oxide  formed,  being  a  base,  would  dissolve  in  sul- 
phuric acid,  forming  copper  sulphate  and  water :  — 

Cu  0  +  H2  S04  =  Cu  S04  +  H2  0, 
so  that  the  entire  reaction  would  be :  - 

2  H2  S04  +  Cu  =  Cu  S04  +  S02  +  H2  0. 

Reactions  of  this  kind  are,  however,  not  so  simple  as  we  are  apt  to 
believe  is  the  case. 

Another  method  quite  frequently  employed  in  the  preparation 
of  sulphurous  acid,  is  by  an  addition  of  an  acid  to  a  sulphite ;  for 
example,  — 

Na,  S03  +  2  H  Cl  =  2  Na  Cl  +  H2  S03 . 

Na2  S08  +  H2  S04  =  Na2  S04  +  H2  S08 , 

the  sulphurous  acid  so  formed  then  breaks  down  into  sulphur 
dioxide  and  water.  This  way  of  preparing  sulphur  dioxide  is  often 
very  convenient,  for  the  same  apparatus  which  is  employed  in 
the  production  of  hydrochloric  acid  can  be  used. 

Sulphur  dioxide  is  a  colorless  gas,  with  the  familiar  odor  of  a 
burning  sulphur  match.  The  application  of  the  moderate  cold  of 
salt  and  snow  changes  it  to  a  clear  liquid,  which  boils  at  —  10°.43 
If  this  liquid  is  evaporated  rapidly  under  the  air  pump,  the  temper- 


138      SULPHUR   DIOXIDE;    PROPERTIES,    SULPHUR    TRIOXIDE. 

ature  sinks  to  —  68°,  while  the  liquid  freezes  at  —  76°.  Sulphur 
dioxide  is  poisonous;  when  it  is  present  in  small  quantities  it 
causes  irritation  of  the  throat  and  violent  coughing ;  in  larger  quan- 
tities, hemorrhages  from  the  lungs,  mouth,  and  nose  occur.  Work- 
men who  are  continually  exposed  to  the  gas,  are  affected  with  loss 
of  appetite  and  headache.  Vegetation  is  destroyed  by  sulphur 
dioxide,  so  that  in  many  places  very  stringent  laws  are  passed 
regulating  the  working  of  factories,  from  the  chimneys  of  which 
sulphur  dioxide  escapes.  Sulphur  dioxide  is  not  combustible,  a  fact 
which  is  self-evident  when  we  consider  that  it  is  the  only  compound 
of  sulphur  ever  formed  by  burning  that  element  in  the  air ;  it  is  an 
extremely  stable  body,  and  hence  will  not  support  combustion ;  in 
fact,  only  in  a  few  instances,  such  as  in  its  action  on  sulphuretted 
hydrogen,  does  it  appear  as  an  oxidizer.  (See  page  92.) 

When  sulphur  burns  in  oxygen,  no  change  of  volume  of  the  gas 
occurs ;  this  phenomenon  is  like  that  which  we  observed  in  the  de- 
composition of  sulphuretted  hydrogen  by  means  of  a  hot  iron,  for 
that  also  took  place  without  alteration  of  volume.  The  explanation 
is  the  same  in  both  cases ;  in  the  one,  each  molecule  of  H2  S 
yields  a  corresponding  molecule  of  hydrogen,  while  the  volume  of 
solid  sulphur  produced  need  not  be  taken  in  consideration ;  in  the 
other  each  molecule  of  oxygen  takes  up  an  atom  of  sulphur  to  pro- 
duce a  molecule  of  S02 ;  so  that  any  number  of  molecules  of  H2  S 
would  yield  the  same  number  of  molecules  of  H2 ,  and  any  number 
of  molecules  of  02  would  yield  the  same  number  of  S02 . 

Sulphur  dioxide  can,  under  proper  conditions,  readily  add 
oxygen  to  form  sulphur  trioxide ;  so,  for  instance,  sulphur  trioxide 
is  produced  by  passing  a  mixture  of  sulphur  dioxide  and  oxygen 
through  a  heated  tube  containing  a  piece  of  platinized  asbestos,  or 
by  exposing  sulphur  dioxide  to  the  action  of  ozone.  The  gas  can 
add  chlorine,  just  as  it  can  oxygen,  for  sulphur  dioxide  and  chlo- 
rine, mixed  and  placed  in  the  sunlight,  produce  sulphuryl  chloride, 
S02C12,  a  compound  which  is  of  considerable  importance  to  us 
theoretically :  -  S02  +  2  Cl  =  SO,  CU . 

Sulphur  dioxide  is  quite  soluble  in  water ;  one  volume  of  water 
absorbs  45  volumes  of  the  gas  at  ordinary  temperatures ;  the  solu- 
tion has  the  odor  and  characteristics  of  sulphur  dioxide,  and  con- 
tains sulphurous  acid.*  Sulphurous  acid  is  a  much  more  reactive 

*  A  hydrated  sulphurous  acid  having  a  definite  crystalline  form  has  been 


SULPHUROUS   ACID;   OXIDATION   OF.  139 

substance,  chemically,  than  the  gaseous  anhydride ;  it  readily  absorbs 
oxygen  from  the  atmosphere,  and  therefore,  if  left  exposed  for  any 
considerable  length  of  time,  will  contain  nothing  but  sulphuric  acid. 
It  is  consequently  self-evident  that  oxidizing  agents,  such  as  chlo- 
rine, bromine,  or  nitric  acid,  will  change  sulphurous  acid  to  sul- 
phuric acid  with  the  greatest  ease.  Two  atoms  of  chlorine  or 
bromine  liberate  one  atom  of  oxygen  from  one  molecule  of  water :  — 

H20+2C1  =  2HC1  +  0, 

and  one  formula  weight  of  sulphurous  acid  requires  one  atom  of 
oxygen  to  change  it  to  sulphuric  acid  :  — 

H2S03  +  0  =  H2S04, 
therefore         H2  S03  +  2  X  +  H2  0  =  H2  S04  +  2  HX, 

where  X  is  used  to  designate  the  halogen.  The  oxidation  of  sul- 
phurous acid  by  means  of  nitric  acid  and  the  oxides  of  nitrogen, 
which  is  used  in  one  of  the  most  important  commercial  processes 
known,  —  namely,  in  the  preparation  of  sulphuric  acid,  —  will  be 
discussed  in  connection  with  that  substance.  Sulphurous  acid  is 
one  of  the  favorite  reducing  agents  in  the  laboratory,  and  we  shall 
frequently  have  occasion  to  refer  to  it  as  such. 

When  heated  in  a  sealed  tube,  sulphurous  acid  changes  to 
sulphuric  acid  and  sulphur  :  — 

3  H2  S03  =  2  H2  S04  +  H2  0  +  S, 

and  in  the  same  way  the  sulphites,  when  heated,  always  form  sul- 
phates by  using  all  of  their  oxygen  for  this  purpose,  for  instance :  — 

4  Na2  S03  =  3  Na2  S04  +  Na2  S. 

These  reactions  remind  us  most  forcibly  of  those  which  take  place 
with  the  oxygen  acids  of  the  halogens,  for  those  acids  and  salts 
of  chlorine,  bromine,  or  iodine  which  contain  the  most  oxygen,  are 
also  the  most  stable. 

Acids  which  contain  one  atom  of  hydrogen  replaceable  by  metals 
are  termed  unibasic ;  those  with  two,  dibasic;  those  with  three, 
tribasic ;  those  with  four,  quadribasic,  and  so  on ;  while  all  acids 
with  more  than  one  replaceable  hydrogen  atom  are  polybasic ;  ac- 
cording to  this  nomenclature  sulphurous  acid  is  a  dibasic  acid.  We 

isolated.  This  hydrate  has  the  formula  H2SO3  4-  6  H2O.  This  hydrate  liber- 
ates sulphur  dioxide  even  at  ordinary  temperatures;  it  is  completely  decom- 
posed at  71°.  Graham-Otto,  vol.  4  [2],  1471. 


140       DIBASIC    ACIDS  ;    PRIMARY   AND   SECONDARY   SALTS. 

are  acquainted  with  two  series  of  salts  derived  from  dibasic  acids, 
accordingly  as  the  metal  replaces  one  or  both  atoms  of  hydrogen 
in  a  formula  weight  of  the  acid.  If  we  designate  a  univalent 
metal  by  M',  then  these  salts  of  sulphurous  acid  would  be  M'  HS03 
and  M'2S03,  respectively;  if  M"  denotes  a  divalent  metal,  then 
the  formulae  are  :  — 

H    -S03 

M"  (          =  M"  ( HS03>  and  M"  S03 . 

H    _1S03 

Salts  formed  by  replacing  one  atom  of  hydrogen  in  a  molecule  of 
an  acid  by  a  metal,  are  called  primary;  those  by  replacing  two, 
secondary ;  those  by  replacing  three,  tertiary,  and  so  on ;  so  that 
NaHS03  would  be  primary,  and  JS"a2S03  secondary  sodium  sulphite.* 
When  a  base  acts  upon  sulphurous  acid  we  can  consider  the  first  re- 
action to  be  as  follows  :  — 

_0  —  H+MOH  (  — 0  —  M+H20 


S  - 


(_ 0— H  (— 0— H 

S03  H2  +  MOH  =  S03  HM  +  H2  0. 

Sulphurous  acid  -+-  a  base    =  A  primary  sulphite  +  water. 
More  of  the  base  acting  on  the  primary  salt  would  then  produce 
the  secondary :  — 

f—0  —  M  (— 0  —  M 

S    -o  -S    -o 

(  —0  — H  +  MOH  (— 0  —  M+H2O. 

S03  MH  +  MOH  =  S03  M2  +  H2  0. 

Primary  sulphite    -f  a  base     =  Secondary  sulphite  +  water. 
Adding  a  base  to  a  primary  salt  will  therefore  produce  a  second- 
ary one;  and  inversely,  adding  more  of  the  acid  to  the  secondary 
salt  will  produce  the  primary :  - 

S03  M2  +  H2  S03  =  2  S03  MH. 

That  these  reactions  must  take  place  is  evident  when  we 
consider  that  the  two  hydrogen  atoms  in  sulphurous  acid  belong  to 
different  hydroxyl  groups,  and  that  they  therefore  are  as  independ- 
ent of  each  other  as  if  they  belonged  to  different  acids. 

Salts  which  contain  a  portion  of  the  replaceable  hydrogen  of  the 

*  Sometimes  called  the  acid  and  the  neutral  sodium  sulphite;  the  primary 
sulphite  is  also  sometimes  termed  sodium  bisulphite. 


RELATIVE   STRENGTH    OF   ACIDS.  141 

acid  are  frequently  termed  acid  salts,  and  those  which  have  ex- 
changed all  of  their  hydrogen,  neutral  salts ;  but  such  designations 
are  frequently  misleading,  for  we  are  acquainted  with  salts  of  the 
former  class,  such  as  the  primary  carbonate  of  sodium,  NaHC03, 
which  have  a  neutral,  or  even  a  slightly  alkaline,  reaction;  while 
in  some  cases  those  of  the  latter  are  acid  toward  litmus,  as  in  the 
case  with  aluminium  sulphate,  A12(S04)3.  Where  a  metal  like 
sodium,  which  has  most  pronounced  metallic  properties,  replaces 
the  hydrogen  of  a  weak  acid,  the  resulting  salt  is  apt  to  have 
an  alkaline  reaction ;  and,  conversely,  where  a  metal  which  is 
not  strongly  characterized,  like  aluminium,  is  found  forming  a 
salt  with  a  strong  acid,  then  the  reaction  of  the  salt  will  probably 
be  acid.  But  what  is  a  weak  and  what  is  a  strong  acid?  The 
question  is  much  easier  to  ask  than  to  answer ;  but  the  following 
must  at  present  be  taken  as  coming  nearest  the  truth.  When  an 
acid  acts  on  a  salt  it  partially  expels  the  acid  combined  in  the  lat- 
ter, and  unites  with  the  base,  and  if  equivalent  quantities  are  taken, 
the  distribution  ratio  of  the  base  is  a  measure  of  the  affinity. 
By  equivalent  quantities  we  mean  the  amount  in  grams  expressed 
by  the  formula  weights  representing  one  atom  of  hydrogen  in  each 
acid.  Thus  equivalent  quantities  of  sodium  chloride  and  sulphuric 
acid  would  be  :  — 

Na  Cl,  formula  weight  23  +  35.5  =  58.5 

iH2S04j      ii  „       2+32+    (4X16)   =49 

The  equivalent  quantities  :  — 

58.5  grams  sodium  chloride  and  49  grams  sulphuric  acid,  for  Na  Cl  and  £  H2  SO4 . 

"  Daily  experience  in  the  laboratory  teaches  us  that  the  affinity 
of  acids  for  bases  is  of  such  a  nature  as  to  appear  a  specific  property 
of  the  acids.  When  we  say  that  carbonic  acid  is  a  weak  acid  and 
sulphuric  acid  a  strong  one,  we  do  not  thereby  mean  that  it  is  so 
with  respect  to  this  or  the  other  base,  but  that  it  is  so  in  general."  * 
Any  metal  replacing  the  hydrogen  of  acetic  acid,  for  instance,  will 
be  more  feebly  united  in  the  resulting  salt  than  it  would  be  with 
sulphuric  acid,  no  matter  if  we  compare  the  salts  of  the  intensely 
metallic  potassium  or  of  the  weakly  metallic  aluminium.  If  we 
mix  sodium  chloride  and  nitric  acid  in  aqueous  solution  and  in 

*  Ostwald,  Outlines  of  General  Chemistry,  p.  337  and  sub. 


142  RELATIVE    STRENGTH    OF   ACIDS. 

equivalent  quantities,  the  following  changes  will  take  place,  until 
an  equilibrium  is  reached  :  — 

NaCl-hHN03  =  NaN08  +  H  Gland 
Na  N03  +  H  Cl     =  Na  Cl     +  HN03 . 

So  that  NaCl,  HN03,  NaN03,  and  HC1  will  all  be  present.  Now, 
the  two  salts,  sodium  chloride  and  sodium  nitrate,  will  be  formed 
in  equal  quantities,  so  that  hydrochloric  and  nitric  acids  are  acids 
of  about  the  same  strength.  If,  on  the  other  hand,  we  mix  sodium 
chloride  with  sulphuric  acid  in  equivalent  quantities,  then  about 
two-thirds  of  the  metal  will  remain  in  sodium  chloride,  while  one- 
third  will  go  to  form  sodium  sulphate,  and  accordingly,  sulphuric 
acid  is  a  weaker  acid  than  hydrochloric  acid,  contrary  to  what 
is  generally  supposed.  If  we  use  molecular  formulae  instead  of 
equivalent  weights,  then  the  reaction  between  sodium  sulphate 
and  hydrochloric  acid  could  be  expressed  by  the  following  equa- 
tion :  — 

3  Xa2  S04  +  6  H  Cl  =  2  H2  S04  +  Na2  S04  +  2  H  Cl  +  4  Na  Cl, 

or,  as  the  primary  sulphate  is  formed  in  the  presence  of  an  excess 
of  sulphuric  acid :  — 

3  Na2S04  +  6  HC1  =  H2S04  +  2  XaHS04  +  2  HC1  +  4  Nad, 

so  that  there  will  be  twice,  as  much  sodium  in  sodium  chloride  after 
the  reaction,  as  there  is  in  sodium  sulphate;  or,  using  the  total 
sodium  as  a  basis  of  calculation,  two-thirds  of  the  sodium  will  go  to 
form  the  chloride,  and  one-third  to  form  the  sulphate.*  The  dis- 
tribution ratio  of  the  metal  used  in  the  above  reactions  would 
therefore  be  as  1:1  with  nitric  and  hydrochloric  acids,  and  as 
2 : 1  with  hydrochloric  and  sulphuric  acids.  This  conclusion,  at 
first  sight,  seems  very  strange  if  we  remember  that  when  toler- 
ably concentrated  sulphuric  acid  is  added  to  sodium  chloride  or 
sodium  nitrate,  sodium  sulphate  is  formed,  and  either  hydro- 
chloric acid  or  nitric  acid  is  given  off,  but  we  must  consider  that, 
after  mixing  salts  and  acids,  the  least  volatile  acid  will,  upon  heat- 
ing, finally  expel  the  more  volatile  ones  from  their  salts,  and,  we 

*  Of  course,  if  no  water  is  present,  hydrochloric  acid,  being  volatile,  will 
pass  off,  .-o  that  no  equilibrium  results  until  all  hydrochloric  acid  has  been 
expelled.  This  equation  is  only  true,  therefore,  where  the  solution  of  hydro- 
chloric acid  in  water  renders  the  latter  not-volatile  (see  pages  57  and  G7). 


RELATIVE   STRENGTH   OF   ACIDS.  143 

must  also  remember,  that  sulphuric  is  much  less  volatile  than  either 
nitric  or  hydrochloric  acid  (see  page  57).  It  is  for  this  reason 
that  sulphuric  acid  is  also  finally  expelled  from  its  salts  when 
they  are  heated  with  much  weaker  not-volatile  acids,  such  as  phos- 
phoric or  silicic.  A  few  of  the  more  important  acids  can  be  writ- 
ten in  the  following  order,  judging  their  strength  by  means  of  the 
relative  amounts  of  metal  which  they  will  take  from  a  chloride, 
when  mixed  with  the  latter  in  equivalent  quantities :  — 

1.  Hydrochloric  acid.  6.  Selenic  acid. 

2.  Nitric  acid.  7.  Phosphoric  acid. 

3.  Hydrobromic  acid.  8.  Hydrofluoric  acid. 

4.  Hydroiodic  acid.  9.  Silicic  acid. 

5.  Sulphuric  acid. 

The  heat  produced  by  neutralizing  an  acid  with  a  base  has  appar- 
ently nothing  to  do  with  the  strength  of  this  acid,  for  the  greatest 
heat  is  produced  by  neutralizing  hydrofluoric  acid,  and  the  next 
greatest  with  sulphuric  acid ;  so  that  if  we  were  to  take  the  rela- 
tive heats  of  neutralization  as  indicating  the  relative  strengths  of 
acids,  quite  a  different  order  from  the  one  given  above  would  re- 
sult, and  certainly  those  salts  which  give  the  greatest  amount  of 
heat  in  their  formation,  would,  in  the  dry  state,  require  the  most 
energy  for  their  decomposition.  At  some  future  time  the  relation 
between  the  heat  of  neutralization  and  the  strength  of  an  acid 
will  undoubtedly  be  discovered. 


144          SULPHUR  TKIOXIDE;   PREPARATION,  PROPERTIES. 


CHAPTER   XXI. 

SULPHUR    TRIOXIDE,    SULPHURIC    ACID,    AND    THE    REMAINING 
SULPHUR  ACIDS. 

Sulphur  trioxide  ;  Formula,  S03 ;  specific  gravity,  1.97  at  20° ;  spe- 
cific gravity  of  vapor,  2.76,  H2  =2,  is  79.48. 

SULPHUR  dioxide  does  not  unite  with  oxygen  under  ordinary 
circumstances,  but  when  a  mixture  of  the  two  gases  is  passed  over 
heated  platinized  asbestos,  union  takes  place  and  sulphur  trioxide 
is  formed.  When  the  latter  is  required  for  use  in  any  quantity  it  is 
better  to  heat  fuming  sulphuric  acid,  as  this  substance  contains  a 
large  quantity  of  the  trioxide  dissolved,  or  else  to  heat  an  easily 
decomposed  sulphate  which  will  form  a  base  and  sulphur  trioxide. 
Ferric  sulphate  is  best  for  this  purpose ;  the  reaction  takes  place  as 
follows :  — 

Fe2(S04)3  =  Fe208  +  3S08. 

The  action  of  ozone  on  sulphur  dioxide  also  produces  the  trioxide 
(page  138). 

Pure  sulphur  trioxide  is  a  colorless  liquid  at  ordinary  tempera- 
tures, but  when  gradually  cooled  it  forms  colorless  prismatic  crystals 
which  melt  at  15°.  The  substance  boils  at  46°,  and  forms  a  color- 
less vapor  which  has  a  specific  gravity  of  2.76,  air  being  one,  and 
therefore,  hydrogen  being  two,  of  79.48 ;  its  molecular  weight  is  con- 
sequently 80,  and  the  formula  S03.  The  ordinary  sulphur  trioxide 
of  commerce  is  a  substance  crystallizing  in  felt-like  crystals  resem- 
bling asbestos,  and  this  form  has  been  taken  for  a  second  modifica- 
tion of  the  body;  in  all  probability,  however,  this  difference  in 
appearance  is  due  to  the  presence  of  traces  of  water.  Pure  sul- 
phuric anhydride  does  not  redden  litmus  paper  nor  does  it  attack 
the  hands ;  the  acid  and  corrosive  properties  only  appear  when,  by 
the  addition  of  water,  the  substance  is  converted  into  sulphuric 
acid.  Sulphur  trioxide  greedily  absorbs  moisture  from  the  atmos- 
phere ;  when  a  little  of  the  substance  is  placed  in  water  it  unites 
with  the  latter  with  a  hissing  noise,  like  that  produced  in  immers- 


SULPHURIC   ACID;   CONSTITUTION.  145 

ing  a  red-hot  iron.     When  heated  to  a  red  heat,  sulphur  trioxide  is 
dissociated,  forming  sulphur  dioxide  and  oxygen  in  the  proportion 
of  two  volumes  of  the  former  to  one  of  the  latter :  — 
=0       0=)  r=00=| 

=o  +  o  =  U   =   s  |  =  o  o  = ;  s. 

=  0      0  =  )  "  0"  ="0" 

Sulphur  trioxide  is  the  anhydride  of  sulphuric  acid  and  yields  the 
latter  on  addition  of  water ;  S03  -f  H2  0  =  H2  S04 . 

The  constitution  of  sulphuric  acid  will  best  be  understood  if  we 
consider  its  formation  from  sulphuryl  chloride,  S02  C12 .  We  have 
seen  that  sulphur  dioxide  and  chlorine  unite  to  form  sulphuryl  chlo- 
ride (page  138).  This  change  can  be  represented  by  structural 
formulae  as  follows :  — 


ci 


=  0 
=  0 


r; 


s 


=  0 

=  0 


Cl  I  — Cl. 

Now,  sulphuryl  chloride,  when  water  is  added,  breaks  down  into 
sulphuric  acid  and  hydrochloric  acid,  and  it  is  obvious  that  in  such 
a  reaction  the  chlorine  atoms  must  be  replaced  by  hydroxyl  groups, 
so  that  sulphuric  acid  must  contain  two  of  the  latter.  The  follow- 
ing formulae  will  make  this  conclusion  apparent :  — 

_C1  +  H  —  0  —  H  r  — 0  —  H  +  HC1 

=  0    '  I   =0 

=0  5 1 =0 

^_  Cl  +  H  —  0  —  H  I  —  0  —  H  -f  H  Cl. 

Sulphuric  acid  can  successively  take  up  one  and  two  molecules 
of  water  to  form  two  hydrated  acids,  H4S05  (H2S04  4-  H20),  and 
H6S06  (H2S04  +  2H20)  ;  it  is  in  the  latter  form  that  the  acid 
probably  exists  when  a  large  excess  of  water  is  present,  so  that  the 
three  oxygen  atoms  of  sulphuric  anhydride,  when  the  latter  is  dis- 
solved, finally  give  place  to  six  hydroxyl  groups :  — 

— O— H 
— O— H 
— O— H 
— O— II 
— O— H 


=0 


S\  =0  =    S     = °+H 

— O — H 

l_O— H 


0-H 
320   =  H.2S04,  +  H20  =  H4S05,  +  H2O          =H0SO6. 


146  SULPHURIC  ACID;  HISTORY. 

The  two  acids,  H2  S04  -f  H2  0,  and  H2  S04  +  2  H20,  are  known, 
but  by  far  the  greater  number  of  sulphates  are  derived  from 
H2S04. 

Sulphuric  acid  is  one  of  the  most  important  commercial  pro- 
ducts, and  the  part  which  it  plays  in  modern  civilization  is 
fundamental.  Until  the  use  of  the  ammonia  process,  it  was  essen- 
tial for  making  soda,  and  soda  is  required  to  produce  both  soap 
and  glass.*  It  is  true  that,  if  we  can  measure  the  civilization 
of  a  nation  by  the  amount  of  soap  which  it  uses,  the  quantity  of 
sulphuric  acid  consumed  can  with  greater  reason  be  taken  as  an 
indication  of  the  stage  of  development  arrived  at  by  a  people. 

The  acid  has  been  known  since  the  time  of  the  Arabian  alche- 
mists, but  it  was  first  accurately  described  by  Basil  Valentine,  who 
prepared  it  in  the  fifteenth  century  by  heating  green  vitriol  (ferrous 
sulphate,  Fe  S04)  with  sand.  The  French  emigrants  of  the  reign  of 
Louis  XIV.  taught  the  English  how  to  prepare  the  acid  by  oxidiz- 
ing sulphur  with  nitre,  and  a  quack  doctor  named  Ward,  using  this 
method,  established  the  first  sulphuric  acid  factory  in  England,  at 
Richmond.  Dr.  Ward  partly  filled  a  glass  vessel  of  about  200 
litres  capacity  with  water,  and  then  placed  within  this  an  earthen 
pot  containing  an  iron  ladle,  on  which  was  burning  a  mixture  of 
sulphur  and  saltpetre,  keeping  the  whole  tightly  covered  until  the 
combustion  was  complete.  Of  course,  this  was  all  very  crude,  and 
the  product  proportionally  dear,  but,  nevertheless,  it  was  a  vast 
improvement  on  the  old  alchemistic  method,  for  the  price  of  sul- 
phuric acid  was  reduced  from  about  three  dollars  and  twenty-five 
cents  to  sixty  cents  a  pound.  In  1746  the  glass  vessel  was  replaced 
by  a  lead  chamber,  and  after  this  improvement  exportation  to  the 
Continent  began,  so  that  the  acid  became  known  as  English  sul- 
phuric acid,  a  name  which  it  bears  to  the  present  day.  The  next 
improvement  consisted  in  substituting  steam  for  water  and  in  burn- 
ing the  mixture  of  saltpetre  and  sulphur  outside  the  chamber, 
while  the  water  vapor  was  passed  through  a  flue  together  with  the 
products  of  oxidation,  the  process  by  this  means  becoming  continu- 
ous. Finally,  the  sulphur  was  burned  in  a  separate  furnace,  while 
the  sulphur  dioxide  so  formed  was  oxidized  in  the  chamber  by 
means  of  nitric  acid;  and  so  the  present  continuous  process  was 
evolved  from  Dr.  Ward's  glass  vessel,  while  the  price  of  the  acid 
sank  from  sixty  cents  to  about  one  cent  a  pound. 

*  See  p.  390. 


SULPHURIC   ACID;   MANUFACTURE.  147 

The  manufacture  of  sulphuric  acid  is  based  upon  the  following 
changes  :  -  S02  +  H2  0  =  Hs  S03  , 

H2S03  +       0  =  H2S04, 

and,  therefore,  its  cheapness  depends  upon  the  ease  with  which  sul- 
phur dioxide  can  be  prepared,  and  upon  the  substance  used  as  an 
oxidizer.  It  might  not  be  unreasonable  to  suppose  that  the  oxygen 
of  the  atmosphere  would  be  most  available  for  the  purpose  ;  but, 
unfortunately,  sulphurous  acid  is  oxidized  much  too  slowly  for  com- 
mercial purposes  by  oxygen  alone.  If,  however,  we  could  furnish 
some  ready  means  of  conveying  oxygen  from  the  atmosphere  to  the 
sulphurous  acid  by  the  intervention  of  some  chemical  compound,  the 
problem  would  be  solved.  This  is  done  by  means  of  nitric  acid  in 
the  present  continuous  process  for  the  manufacture  of  sulphuric 
acid,  in  which  the  main  changes  taking  place  are  as  follows  :  — 

1.  In  a  mixture  of  nitric  acid  and  sulphurous  acid,  each  com- 
ponent acts  on  the  other  ;  the  sulphurous  acid  is  oxidized  to  sul- 
phuric acid,  and  the  nitric  acid  is  reduced  to  nitrous  acid. 


Nitrous  acid,  like  other  acids  the  anhydrides  of  which  are  gases, 
at  once  breaks  down  into  water  and  nitrous  anhydride  :  — 


so  that  the  entire  change  can  be  represented  by  the  equation  :  — 


2.  Nitrogen  trioxide  (N2  03)  with  water,  sulphur  dioxide,  and 
oxygen  forms  a  compound  known  as  nitrosyl-sulphuric  acid,  which 
latter  is  simply  sulphuric  acid  in  which  an  hydroxyl  group  is  sub- 
stituted by  the  group  of  elements  N02  .* 

f  —OH 

n   I   =0  n 

Ol  =0  0 

L—  OH. 

Sulphuric  acid.  Nitrosyl-sulphuric  acid. 

*  The  group  NO2  is  called  the  nitro  group  ;  it  is  univalent,  just  as  is  the 
hydroxyl  group.  The  formula  of  nitrosyl-sulphuric  acid  may  possibly  be  re- 

f—  O—  N  =  0 
presented  by  JO  ^  o  .      This  would  be  sulphuric  acid,  in  which   one 

[—  0-H 
hydrogen  atom  has  been  replaced  by  the  univalent  group,  —  N  =  O  (the  nitroso 


148  SULPHUKIC    ACID'    MANUFACTURE. 

The  reaction  can  be  represented  as  follows :  — 

(OH 

2  S02  +  N2  03  +  2  0  +  H2  0  =  2      Q  J  02 

°(N02. 

3.  Nitrosyl  sulphuric  acid  breaks  down  into  X2  03  and  sulphuric 
acid  on  addition  of  water  :  — 

fOH  (OH 

a.  S  \  02  =      S  ]  O2 

1  :-•- :  (OH  +  N02H, 

U   N02  +  H    I  OH 

b.  2  N02  H  =  N2  08  +  H2  0,    so   that,  combining   a   and   b}   we 
have : — 

OH 

Oo         +  H20=2H2S04  +  N203. 

NO, 

4.  The  nitrogen  trioxide,  so  regenerated;  can,  with  steam  and 
air,  once  more  form  nitrosyl  sulphuric  acid,  which,  with  water,  will 
form  sulphuric  acid ;  so  that,  theoretically,  an  infinitely  small  quan- 
tity of  nitric  acid,  introduced  at  the  beginning  of  the  operation, 
would  oxidize  any  amount  of  sulphur  dioxide.     That  this  is  not 
in  reality  the  case  is  due  to  the  fact  that  other  minor  reactions, 
producing  lower  oxides  of  nitrogen,  take  place ;  and  also  because 
the  nitrogen  of  the  air,  as  it  takes  no  part  in  the  reaction,  gradually 
dilutes  the  gases  to  such  an  extent  as  to  render  loss  inevitable. 
The  commercial  production  of  sulphuric  acid  is  carried  on  in  works 
a  simple  diagram  of  which  is  shown  in  Fig.  9.44     Sulphur  dioxide  is 
prepared  by  burning  either   sulphur  or  iron  pyrites   (FeS2)   in  a 
furnace  with  free  access  of  air ;   the  gas  enters  the  flue  (V),  and  is 
conducted  to  the  top  of  the  tower  (G)  which  is  filled  with  pieces  of 
fire-brick.     Two  vats  (&),  one  containing  dilute  sulphuric  acid,  and 
the  other  a  concentrated  acid  in  which  oxides  of  nitrogen  are  dis- 
solved, are  constantly  emptying  their  contents  into  G.     The  concen- 
trated acid  is  supplied  from  the  tower  (G'),  the  object  of  which  we 
shall  see  later  on.     Concentrated  sulphuric  acid  can  dissolve  large 
quantities  of   the  oxides  of  nitrogen,  but  dilute  acid  has  no  such 
power,  so  that    mixing  the  contents  of   the  two  receptacles  at  b 

group).  This  interpretation  is  rendered  probable  by  the  ease  with  which 
nitrosyl-sulphuric  acid  is  broken  down  by  water,  as  bodies  containing  the  nitro 
group  are,  as  a  rule,  more  stable. 


SULPHURIC    ACID;    MANUFACTURE. 


149 


liberates  these  oxygen  compounds.  The  acid  and  the  oxides  of 
nitrogen  mingle  with  the  sulphur  dioxide  entering  at  a ;  the  hot 
gas  serves  to  concentrate  the  dilute  acid,  while  the  latter  cools  the 
gas  before  it  passes  into  the  leaden  chambers.  Sulphur  dioxide, 
mixed  with  oxides  of  nitrogen,  now  enters  the  bottom  of  the  chamber 
1,  and,  in  this,  comes  in  contact  with  steam  and  the  vapors  of  nitric 
acid  (the  latter  prepared  by  heating  a  mixture  of  sulphuric  acid  and 
sodium  nitrate)  ;  the  fuel  being  the  sulphur  burning  to  form  sulphur 
dioxide.  Sulphuric  acid  is  formed  in  the  first  lead  chamber,  while 


Fig.  9. 

the  unused  gases  are  passed  into  2,  and  then  into  3,  in  both  of 
which  places  they  come  in  contact  with  more  steam,  so  that,  in 
these,  the  changes  are  completed.  The  air,  which  is  supplied  by 
the  draught  of  a  large  chimney,  gives  up  its  oxygen  in  going  through 
the  chambers,  and  therefore  the  gases  become  so  diluted  with  nitro- 
gen .  as  to  be  no  longer  capable  of  taking  part  in  the  reactions. 
These  diluted  portions  are  passed  irr  at  the  top  of  the  tower  G', 
which  contains  pieces  of  coke,  over  which  concentrated  sulphuric 
acid  is  constantly  trickling ;  here  the  acid  dissolves  the  remaining 
oxides  of  nitrogen  which  pass  from  chamber  3,  and,  saturated  with 
these  gases,  it  can  be  pumped  to  one  of  the  vats  b  above  the  first 
tower,  to  be  used  as  was  indicated  above.  Experience  has  shown 
that  the  acid  collected  on  the  floors  of  the  chambers  must  not  con- 
tain more  than  65-66  per  cent  of  H2  S04 ,  and  must  not  have  a 
specific  gravity  higher  than  1.5 ;  but  as  the  commercial  acid  has  a 


150  SULPHURIC   ACID;   PROPERTIES. 

specific  gravity  of  1.83,  and  must  contain  89-90  per  cent  of  H2  S04 , 
the  concentration  of  £he  chamber  acid  is  carried  farther  by  placing 
it  in  flat  lead  pans  and  evaporating  the  excess  of  water  until  a 
specific  gravity  of  1.75  is  reached ;  and  then,  because  a  stronger 
acid  attacks  lead,  by  finally  completing  the  evaporation  in  platinum 
or  glass  vessels. 

Commercial  sulphuric  acid  is  a  liquid  *  which,  on  superficial  exam- 
ination, appears  like  an  oil ;  it  is  usually  light  brown  in  color,  owing 
to  impurities.  The  latter  consist  of  sulphate  of  lead  (which  is  in- 
troduced from  the  lead  chambers,  and  which  is  never  absent),  of  the 
oxides  of  nitrogen  ( N2  03  and  N02),  of  hydrochloric  acid,  and  of 
arsenious  oxide  (As203)  (the  latter  is  present  because  arsenic  is 
found  in  the  minerals  roasted  for  the  preparation  of  sulphur  diox- 
ide). In  burning  iron  pyrites,  the  iron  is  changed  to  its  oxide,  and 
the  sulphur  to  the  dioxide ;  and,  as  the  pyrites  frequently  contain 
selenium,  selenium  dioxide  and  selenium  collect  in  the  flues  and  in 
the  muddy  residue  at  the  bottom  of  the  chambers.  The  brownish 
color  of  commercial  sulphuric  acid  is  caused  by  organic  substances 
(dust,  etc.)  which  have  fallen  into  it.  Pure  sulphuric  acid  is  pre- 
pared from  the  commercial  article  by  distilling  from  platinum 
vessels.  Many  so-called  pure  acids,  however,  contain  arsenic, 
because  arsenious  oxide  is  volatile,  and  will  therefore  pass  over  in 
the  distillation,  unless  care  has  been  taken  to  previously  oxidize  it 
to  arsenic  acid. 

The  pure  acid,  H2  S04 ,  is  a  colorless,  oily  liquid,  with  a  specific 
gravity  of  1.85  at  0°.  Upon  being  cooled  to  0°,  it  crystallizes  in 
large  prismatic  crystals,  which  melt  at  10.°  5 ;  it  boils  at  338°,  but, 
before  that  temperature  is  reached,  the  acid  begins  to  decompose 
into  sulphur  trioxide  and  water ;  this  separation  is  perfect  if  heated 
somewhat  above  its  boiling  point :  — 

H2S04  =  S03+H20. 

Sulphuric  acid  has  a  great  inclination  to  take  up  water,  and  in 
so  doing  can  form  two  hydrated  acids :  H4  S05  and  H6  S06 .  The 
first  one  of  these  is  formed  when  a  mixture  of  sulphuric  acid,  with 
just  enough  water,  is  cooled  to  a  low  temperature,  then  prisms  of 
H4  S05  separate.  The  second,  or  normal,  hydrate  is  produced  by 

*  The  name  "  oil  of  vitriol "  was  given  to  sulphuric  acid  by  the  alche- 
mists, because  of  its  oily  appearance,  and  because  it  was  first  prepared  from 
green  vitriol  (ferrous  sulphate). 


SULPHURIC   ACID  ;   REACTIONS.  151 

adding  the  requisite  quantity  of  water  to  H4  S05 .  A  large  amount 
of  heat  is  developed  in  the  formation  of  these  hydrates,  yet  the 
heat  production  does  not  «^ase  Avhen  exactly  enough  water  to  pro- 
duce the  normal  hydrate  ha^  been  added ;  it  will  continue  until  the 
proportions  are  expressed  by  H2  S04  +  1600  H2  0,  when  178  K  will 
have  been  developed.  Sulphuric  acid  has  such  a  strong  tendency 
to  unite  with  water  that  it  can  take  the  elements  of  that  compound 
from  organic  substances.  If  it  is  mixed  with  sugar,  starch,  pieces 
of  wood,  or  similar  substances  which  contain  hydrogen  and  oxygen 
in  exactly  the  proportions  to  form  water,  it  will  char  them  as  if 
they  had  been  burned  in  an  insufficient  supply  of  oxygen;  for, 
after  the  hydrogen  and  oxygen  have  been  taken  from  such  bodies, 
nothing  but  carbon  remains.  In  a  similar  way  sulphuric  acid  will 
attack  the  skin  or  mucous  membrane,  so  that  the  concentrated  acid 
is  a  violent  poison. 

Reducing  agents  readily  change  sulphuric  acid  to  sulphur 
dioxide,  and  even  to  sulphur,  or  sulphuretted  hydrogen.  We  have 
studied  examples  of  such  reduction  in  the  changes  which  take  place 
when  hydroiodic  or  hydrobromic  acid  acts  upon  sulphuric  acid  (see 
pages  80  and  85).  As  a  general  rule,  hydrogen  compounds  will 
readily  reduce  sulphuric  acid  if,  like  hydroiodic  and  hydrobromic 
acids,  they  are  unstable ;  for  instance,  hydrogen  sulphide,  selenide, 
or  telluride  will  act  in  the  following  way :  — 

H2S04  +  H2S  =  2  H20  +  S02  +  S. 

This  reaction  is  exactly  like  that  taking  place  between  hydro- 
bromic and  sulphuric  acids  :  — 

H2  S04  +  2  H  Br  =  2  H2  0  +  S02  +  2  Br, 

only  in  the  one  case  sulphur,  and  in  the  other  bromine,  is  produced. 
We  have  already  discussed  the  reduction  of  concentrated  sul- 
phuric acid  by  metals  (page  136),  so  that  in  this  place  it  is  only 
necessary  to  add  that,  in  addition  to  copper  or  zinc,  silver,  mercury, 
and  a  number  of  other  metals,  will  produce  sulphur  dioxide  when 
they  are  heated  with  sulphuric  acid ;  but  we  must  remember  that, 
in  cases  where  metals  are  attacked  by  the  dilute  acid,  hydrogen  is 
.liberated,  as  we  saw  when  we  discussed  the  preparation  of  that  ele- 
ment (see  page  31).  Other  easily  oxidized  substances,  such  as  char- 
coal and  sulphur,  will  also  readily  reduce  sulphuric  acid.  Diluted 
sulphuric  acid  must  possess  much  less  chemical  energy  than  does 


152  SULPHATES;   PRIMARY  AND   SECONDARY. 

the  concentrated  liquid ;  as  a  consequence,  it  will  take  more  energy 
to  decompose  it.  This  fact  is  evident  when  we  recall  the  great 
amount  of  heat  which  is  given  off  when  sulphuric  acid  is  dissolved 
in  water.  It  is  because  of  its  relative  instability  that  concentrated 
sulphuric  acid  is  so  easily  reduced. 

One  of  the  chief  laboratory  uses  for  sulphuric  acid  is  to  prepare 
other  acids  by  its  action  on  the  salts  of  the  latter,  and  we  have 
already  encountered  a  number  of  cases  in  which  the  acid  was  used 
for  this  purpose  ;  for  example  :  — 

2KN03     +H2S04  =  K2S04    +  2HN03, 
2  K  Cl  04  +  H2  S04  =  K2  S04    +  2  H  Cl  04 , 
Na2  S03  +  H2  S04  =  Na2  S04  +     H2  S03 , 

2  Na  Cl     +  H2  S04  =  Na2  S04  +  2  H  Cl. 

Heretofore,  in  studying  such  reactions,  we  have  always  taken 
the  formation  of  the  secondary  sulphate  for  granted ;  but  this,  in 
reality,  does  not  take  place  if  an  excess  of  sulphuric  acid  is  present. 
Sulphuric  acid  is  dibasic,  and  can  therefore  form  two  series  of  salts, 
the  primary  (MH  S04 ),  and  the  secondary  (M2  S04 ).  Now,  if  we 
compare  the  action  of  sulphuric  acid  on  sodium  nitrate  or  sodium 
chloride  with  that  of  the  same  acid  on  sodium  hydroxide,  we  shall 
see  that  they  are  analogous  processes ;  and,  if  we  can,  in  acting  on 
sulphuric  acid  with  sodium  hydroxide,  replace  first  one  and  then 
both  hydrogen  atoms  with  the  metal,  it  follows  that  we  should  have 
the  same  changes  were  we  to  substitute  sodium  chloride  or  sodium 
nitrate :  — 

1.  XaOH  +  H— O^O  =O_H9O  =Xa—  O  ^0  =O 

H— O^O=O  H  —0^0=0 

Xa  OH  +  H2  SO4  =  H.2  O  +  Xa  H  SO4 

2.  Xa— O.   0  =O=  Xa— O..O=O 
XaOH  +  H— O^O=O     H2O  +  Xa—  O^O=O 

Xa  OH  +  Xa  HSO4  =  H2  O  +  Xa2  SO4 

1.  XaCl    +  H— O.    0=0_HC1  +Xa— O.   0=0 

H— O>0=O  H   —0^0=0 

Xa  Cl    +  H2  SO4  =  H  Cl  +  Xa  HSO4 

2.  Xa— O.   0=O=  Xa— O^O=O 
XaCl    +  H—  0-^0=0      HC1 +Xa— 0^0=0 

Xa  Cl    +  Xa  HS04  =  H  Cl  +  Xa2  SO4 . 

The  above  reactions  are  a  necessary  result  of  the  fact  that,  in 
polybasic  acids,  the  hydrogen  atoms  are  entirely  independent  of 


SULPHATES;  PRIMARY  AND  SECONDARY.  153 

each  other.  If,  therefore,  in  decomposing  a  salt,  we  use  an  excess 
of  sulphuric  acid,  the  primary  sulphate  results ;  if  we  use  an  excess 
of  the  salt,  we  produce  the  secondary  sulphate,  for,  comparing  the 
two  reactions :  — 

NaCl  +  H2S04  =  NaHS04  +  H  Cl,  and  2  NaCl  +  H2SO4  =  Na2SO4  +  2  HC1, 

we  see  that,  in  the  second  one,  we  have  twice  as  much  sodium  chlo- 
ride in  proportion  to  the  acid  as  in  the  first.  In  laboratory  practice 
it  is  expedient  to  calculate  the  relative  quantities  of  salt  and  sul- 
phuric acid  so  as  always  to  produce  the  primary  sulphate,  because 
it  is  easier  to  fuse  and  more  convenient  to  handle  than  the  secon- 
dary. On  heating  a  primary  sulphate  *  we  form  the  corresponding 
secondary  sulphate  thus  :  — 

2  Na  HS04  =  Na2  S04  +  H2  S04 , 
and  we  can  accomplish  the  same  result  by  adding  a  base :  — 

Na  HS04  +  Na  OH  =  Na2  S04  +  H2  0. 

It  is  obvious  that  the  hydroxide  so  added  may  contain  a  different 
metal  from  that  already  present  in  the  salt,  so  that  secondary  salts 
containing  two  metals  may  be  formed :  — 

Na  HS04  +KOH  =  Na  KS04  +  H2  0. 

By  adding  sulphuric  acid  to  the  secondary  sulphate  we  can  produce 
the  primary  :  — 

Na2  S04  +  H2  S04  =  2  Na  HS04 . 

No  salts  formed  by  replacing  all  of  the  hydrogen  atoms  in 
either  of  the  two  hydrated  acids,  H4  S05  and  H6  S06 ,  exist,  but 
some  are  known  in  which  two  of  these  have  been  substituted  by 
a  metal ;  such  salts  are  frequently  considered  as  being  ordinary  sul- 
phates, with  the  additional  water  attached  in  some  mysterious  way 
known  as  "molecular  union,"  and  so  their  formulae  are  written 
M"S04  4-  H2  0,  and  M"S04  +  2  H2  0  ;  but  it  is  more  rational  to  look 
upon  these  as  secondary  salts,  M"H2  S05  and  M"H4  S06  of  the  hy- 
drated acids,  H4  S05  and  H6  SOG .  This  theory  is  borne  out  by  the 
fact  that  many  of  these  salts  lose  water  only  at  temperatures  con- 
siderably above  that  necessary  to  expel  water  of  crystallization, 
which  fact  seems  to  indicate  that  water,  as  such,  is  not  present  in 
them,  but  that  it  is  in  the  form  of  hydroxyl  groups.  A  number  of 
more  complicated  salts  are  supposed  to  be  derived  from  normal  sul- 

*  Also  termed  an  acid  or  a  bisulphate. 


154  DISTJLPHURIC   ACID;   THIOSULPHURIC   ACID. 

phuric  acid,  but  for  information  regarding  these  a  larger  manual 
must  be  consulted. 

Sulphur  trioxide  is  very  soluble  in  sulphuric  acid  ;  the  solution 
is  a  liquid,  having  an  oil-like  appearance  which  gives  off  dense 
white  fumes  when  it  is  exposed  to  the  air  :  — 


This  substance,  which  has  a  composition  represented  by  the  for- 
mula H2  S2  07  ,  is  termed  fuming  sulphuric  acid  ;  it  is  derived  from 
sulphuric  acid  by  separating  water  from  hydroxyl  groups  belonging 
to  different  molecules,  so  that  its  constitution  would  be  represented 
as  follows  :  — 

r  OH  HO  ]  r  OH  HO  i 

nO  0  In  n      0  On 

o  j  o  .....................     o  ro  Oio  oro 

_io_  H   H|—  oj  -o- 

This  acid,  which  is  termed  disulphuric  or  pyrosulpliuric  acid, 
is,  therefore,  dibasic  ;  it  is  formed  by  linking  two  monovalent 
groups,  S03'H,  by  means  of  a  divalent  oxygen  atom,  and  its  name, 
disulphuric  acid,  suggests  this  constitution.  The  union  of  two 
such  monovalent  groups  by  means  of  a  polyvalent  atom  is  a  phe- 
nomenon of  quite  common  occurrence.  On  adding  water  to  disul- 
phuric acid,  sulphuric  acid  is  formed,  and  on  extracting  water  from 
disulphuric  acid,  sulphur  trioxide  remains;  so  that  this  acid  lies 
between  sulphuric  acid  and  its  anhydride,  bearing  the  same  rela- 
tionship to  sulphuric  acid  as  the  latter  does  to  H4  S05  :  — 

H2S207  +  H20  =  2  H2S04, 

H2S207-H20  =  2S03. 

We  have  seen  that,  because  of  the  great  chemical  similarity 
between  the  two  elements,  sulphur  can  take  the  place  of  oxygen 
in  many  acids.  We  are  acquainted  with  the  salts  of  one  acid 
(tliiosulphurie  acid),  derived  from  sulphuric  acid  by  means  of  such 
a  substitution  ;  the  acid  itself  is  not  known.  Thiosulphate  of  so- 
dium, Na2  S2  03  ,  the  most  common  salt  of  this  acid,  can  be  consid- 
ered as  sulphate  of  sodium,  in  which  one  atom  of  oxygen  has  been 
replaced  by  one  of  sulphur  :  — 

Na2S04,  sodium  sulphate,  and  Na2SS08,  sodium  thiosulphate. 

This  compound  is  frequently  called  the  hyposulphite  of  sodium  ; 


THIONIC   ACIDS. 


155 


but,  obviously,  such  a  name  is  not  advisable,  because  it  suggests  a 
relationship  to  sulphurous  acid  similar  to  that  sustained  by  hypo- 
chlorous  acid  to  chlorous  acid,  while  such  a  parallelism  does  not  in 
reality  exist.  Thiosulphates  are  changed  to  the  sulphates  by  heat- 
ing, all  of  the  oxygen  present  in  the  thiosulphate  being  used  to  form 
the  sulphate,  while  the  excess  of  sulphur  unites  with  the  excess  of 
metal  to  form  the  sulphide.* 


03  =  3 


Na  S 


When  thiosulphuric  acid  is  liberated  from  its  salts  by  the  addition 
of  other  acids,  it  at  once  breaks  down  into  water,  sulphur  dioxide, 
and  sulphur,  f 

Na2  S2  03  +  H2  S04  =  Na2  S04  +  H2  0  +  S02  +  S. 

In  addition  to  the  ones  which  have  been  mentioned,  a  series  of 
acids  which  contain  sulphur,  and  which  have  the  following  for- 
mulae, exists  :  — 


1.   Dithionic  acid,        H-O-S 


=  0    0=) 

=  O     O  =  \  S-O-H  =  H2S2O6 


,=0     0=) 

2.  Trithionic  acid,       H-O-S    \  =  O    O  =  >  S-O-H  =  H2S3O6 


=  0     0  = 


3.   Tetrathionic  acid,  H-O-S    <  =  O     O  =  >  S-O-H  =  H2S4O( 


,  =0  0  = 

4.   Pentathionic  acid,  H-O-S    ]  =  O  O  =  J>  S-O-H  =  H2S5O6. 

'  — S  — S  — S- 

*  The  formulae  of  the  sulphides  of  some  metals,  notably  those  of  the 
alkali  metals,  certainly  bear  a  most  remarkable  resemblance  to  the  oxygen 
compounds  we  have  just  been  studying.  Thus  we  have  sulphides  of  potassium, 
K2  S2,  Ko  S8,  K2  S4,  and  K2  S5,  called  polysulphides,  the  last  two  of  which  when 
written  K2  SS8  and  K2  SS4  might  possibly  be  K2  SO3  and  K2  SO4,  in  which  oxy- 
gen is  replaced  by  sulphur  ;  as,  however,  the  parallelism  does  not  extend 
beyond  the  mere  relationship  in  the  number  of  atoms,  and  as  we  have  no 
knowledge  of  the  structural  formulae  of  the  polysulphides,  this  interpretation 
is  purely  speculative.  (See  Drechsel;  Journal  fur  Praktische  Chemie, 
[2]  4,  20.) 

t  The  liberation  of  sulphur  on  addition  of  acids  distinguishes  thiosul- 
phates  from  sulphites. 


156  SULPHUR;  HALIDES  OF. 

These  acids,  in  all  probability,  contain  the  univalent  group  of 
elements  :  — 

(  =0 

-SJ-o 

(_0  —  H, 

which  also  occurs  in  disulphuric  acid  (see  page  154).  Two  of 
these  groups  are  united  in  dithionic  acid,  while  in  the  remaining 
three  they  are  joined  by  means  of  sulphur  atoms,  as  is  shown  by 
the  formulae.  Trithionic  acid,  therefore,  is  disulphuric  acid  in  which 
the  linking  oxygen  atom  is  replaced  by  one  of  sulphur.  A  larger 
text-book  must  be  consulted  for  the  methods  of  preparation  and 
general  characteristics  of  these  acids.  The  constitutional  formulae 
of  the  compounds  which  have  been  discussed  in  that  part  of  the 
work  following  sulphuric  acid,  express  the  present  state  of  our 
knowledge,  but  the  whole  subject  will  bear  further  investigation. 
Constitutional  formulae  are  constructed  with  the  view  of  so  arran- 
ging the  grouping  of  the  atoms  in  them  that  all  of  the  reactions 
entered  into  by  the  substances  which  they  represent  can  be  ex- 
plained by  them,  and  further,  they  also  frequently  indicate  the  man- 
ner in  which  these  substances  are  formed.  As  soon  as  any  facts 
-contradicting  a  structural  formula  in  general  use  are  discovered, 
the  formula  must  either  be  abandoned  or,  at  least,  so  modified-  as  to 
agree  with  the  new  discovery. 

As  might  be  expected,  sulphur  can  form  unstable  chlorides  by 
direct  union  with  chlorine.  The  first  product  of  the  action  of 
chlorine  on  sulphur  is  the  inonochloride,  S2C12,  if  the  action  of 
•chlorine  is  continued,  the  dichloride,  SC12,  is  formed,  and  lastly, 
if  a  large  excess  of  chlorine  acts  on  the  dichloride  at  a  temperature 
of  —  22°,  the  tetrachloride,  S  C14 ,  results.*  This  latter  compound 
decomposes,  when  warmed  above  —  22°,  liberating  chlorine,  while 
the  dichloride  is  only  stable  below  6°  to  10°.  Corresponding  com- 
pounds, S2Br2  and  S2I2,  exist ;  and  an  iodide,  S  Ie ,  is  also  described.! 

Compounds  of  sulphur  which  contain  both  chlorine  and  oxygen 
are  derived  from  sulphurous  and  sulphuric  acids,  by  substituting 
chlorine  for  hydroxyl  groups.  They  are  termed  acid  chlorides,  and 

*  Compare  the  behavior  of  this  chloride  with  the  unstable  chloride  of 
manganese  (MnCl4),  which  is  supposed  to  be  formed  when  hydrochloric  acid 
acts  on  manganese  dioxide. 

t  Lamers;  Jour,  fur  Prakt.  Chem.;  84,  349. 


SULPHUR;  ACID  CHLORIDES  OF.  157 

are  SOC12,  thionyl  chloride;  S03C1H,  sulphury  1-hydroxyl  chloride 
(chlor-sulphonic  acid);  and  S02C12,  sulphuryl  chloride. 

_0  — H         r—Cl 


—  0  — H 

Sulphurous  acid.  Thionyl  chloride.  Sulphuric  acid.       Chlor-sulphonic      Sulphuryl 

acid.  chloride. 

Following  the  law  which  we  found  to  be  general  with  the  chlorides 
of  the  not-metals,  these  compounds  are  decomposed  into  the  corre- 
sponding acids  by  addition  of  water ;  thionyl  chloride  forming  sul- 
phurous acid,  and  the  last  two  both  yielding  sulphuric  acid.  The 
decomposition  of  sulphuryl  chloride  in  this  way  not  only  has  an 
important  theoretic  bearing  in  the  constitution  of  sulphuric  acid, 
(see  page  138),  but  what  is  more,  the  formation  of  S03  Cl  H  from 
hydrochloric  acid  and  sulphuric  anhydride  :  — 

=  0 

=  0 

=  0  +  HC1  (Z&~* 

shows  a  resemblance  between  hydrochloric  acid  and  water  in  chemi- 
cal behavior ;  for  the  following  reaction  ( 2 )  is  clearly  analogous  to 
reaction  1 :  — 


c\ 


The  existence  of  these  compounds  illustrates  the  similarity  be- 
tween hydroxides  of  the  metals  and  of  the  not-metals,  for  in  both 
classes  of  compounds  the  hydroxyl  groups  can  be  replaced  by  chlo- 
rine ;  though  with  metals  this  substitution  is  much  more  easily 
brought  about  than  with  not-metals.  The  hydroxide  of  the  metal 
has  only  to  be  treated  with  hydrochloric  acid  in  order  to  form  the 
very  stable  chloride  :  — 

K-0-H  +  HC1  =  KC1  +  H-0-H, 

while  in  forming  the  acid  chlorides  some  roundabout  method,  which 
excludes  the  presence  of  water,  must  be  resorted  to  ;  for  these  com- 
pounds are  all  decomposed  by  the  latter  substance. 


158  SULPHUROUS  ACID;  STRUCTURE  OF. 

Chemists   have   lately   believed   sulphurous   acid   to   have  the 
constitution :  — 

r  =  0  r— 0  —  H 


I.     H  —  R-<=0  and  not         II. 

U(_0  —  H, 


-0  —  H, 

so  that,  if  this  view  is  correct,  an  analogy  between  that  acid  and 
thionyl  chloride  does  not  exist.  This  interpretation,  which  has 
been  given  to  some  experiments  made  with  compounds  belonging 
in  organic  chemistry,  seems  to  be  unnecessary  ;  it  is  more  probable 
that  the  replaceable  hydrogen  in  all  of  the  oxy-acids  which  we 
have  studied  is  present  in  the  hydroxyl  groups,  so  that  the  for- 
mula of  sulphurous  acid  would  be  as  is  shown  above.  (II.) 

Three  other  oxides  of  sulphur,  S2  03  ,  S2  07  ,  and  S04  ,  are 
known.  The  first  two  very  readily  decompose,  the  former  into 
sulphur  and  sulphur  dioxide  (2  S203  =  S  -f  3  S02),  and  S2  07  into 
sulphur  trioxide  and  oxygen  (S207  =  2  S03  -f-  0). 

The  following  table  will  make  the  relationship  between  the  sul- 
phur acids  more  apparent  :  — 


OXY-ACIDS    OF    SULPHUR  ;    TABLE    OF. 


159 


§.0000 


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160  SELENIUM  DIOXIDE;  TELLURIUM  DIOXIDE. 


CHAPTER   XXII. 

THE   COMPOUNDS   OF   SELENIUM   AND  TELLURIUM   WITH 
OXYGEN,    AND   WITH   OXYGEN   AND  HYDROGEN. 

SELENIUM  and  tellurium  do  not  manifest  so  great  a  variety  in 
the  formation  of  oxides  and  acids  as  sulphur  5  but  such  oxides  as 
they  do  form  are  constructed  in  a  manner  analogous  to  the  structure 
of  the  compounds  of  the  latter  element.  They  each,  on  burning, 
yield  the  corresponding  dioxide  :  — 

Se  +  2  0  =  Se  02,  and  Te  +  2  0  =  Te  0,, 

and  both  the  dioxides  are  solid  bodies ;  in  fact,  all  of  the  oxides  of 
the  not-metals  having  higher  atomic  weights  are  solids.  Sulphur 
dioxide,  in  solution,  is  a  powerful  reducing  agent,  always  showing 
the  greatest  tendency  to  take  up  oxygen  so  as  to  form  sulphuric 
acid,  while,  on  the  other  hand,  the  dioxides  of  selenium  and  tellu- 
rium exhibit  a  behavior  which  is  exactly  the  reverse ;  they  part  with 
their  oxygen  so  readily  that  even  the  particles  of  dust  which  may 
come  in  contact  with  them  serve  as  reducing  agents,  and  so  liberate 
selenium  or  tellurium.  Of  course,  sulphurous  acid  readily  reduces 
the  two  oxides,  while  it  is  changed  to  sulphuric  acid :  — 

2H2S03+  Se02  =  2H2S04-f  Se, 
2  H2  S03  +  Te  02  =  2  H2  S04  -f  Te, 

and,  as  a  consequence,  selenium,  and  not  selenium  dioxide,  is  found 
in  the  lead  chambers  of  sulphuric  acid  works.  (See  page  102.) 

Selenium  dioxide  is  a  white,  snow-like  solid,  which  is  best  pre- 
pared by  oxidizing  selenium  with  nitric  acid.  It  does  not  melt ; 
but,  on  heating,  it  changes  directly  from  a  solid  to  a  gas,  at  a 
temperature  below  the  boiling  point  of  sulphuric  acid.  Selenium 
dioxide  greedily  absorbs  moisture  from  the  air,  forming  selenious 
acid.  The  latter  is  a  dibasic  acid,  giving  primary  and  secondary 
salts,  MH  Se  03  and  M2  Se  03 . 

Tellurium  dioxide  is  produced  by  the  oxidation  of  tellurium 
with  nitric  acid.  It  is  a  colorless,  crystalline  solid,  which  is  but 
little  soluble  in  water.  As  tellurium  has  such  a  high  atomic 


SELENIC    ACID;    TELLURIC    ACID.  161 

•• 

weight,  its  lower  oxides  begin  to  bear  some  resemblance  to  the 
oxides  of  the  less  pronounced  metals ;  the  dioxide  has  therefore 
scarcely  any  tendency  to  unite  with  water,  nor  is  it  dissolved  by 
weak  bases ;  an  additional  sign  that  tellurium  is  approaching  the 
metals  in  its  character  is  the  fact  that  the  tetrachloride,  Te  C14 ,  is 
not  completely  decomposed  by  cold  water.  Tellurous  acid  is  a 
white  powder  which,  when  warmed,  readily  decomposes  into  tellu- 
rium dioxide  and  water.  It  is  a  dibasic  acid,  and  forms  primary 
and  secondary  salts,  MH  Te  03  and  M2Te  03 . 

Selenium  trioxide  has  never  been  prepared,  but  the  correspond- 
ing acid,  Ho  Se  04 ,  selenic  acid,  can  be  formed  by  oxidizing  selenious 
acid  with  chlorine  or  bromine,  just  as  the  corresponding  sulphur 
compound  is  oxidized  by  the  same  elements.  Solutions  of  selenic 
acid  cannot  be  concentrated  beyond  the  point  at  which  they  contain 
95  per  cent  of  H2  Se  04 ,  because,  after  this  percentage  is  reached, 
the  acid  breaks  down  into  selenious  acid  and  oxygen.  The  rule 
with  the  selenium  acids  is  therefore  exactly  the  reverse  of  that 
which  held  true  as  regards  the  halogen  and  the  sulphur  compounds, 
for  in  dealing  with  those,  the  acids  containing  the  most  oxygen  are 
the  final  products  formed  by  heating  the  ones  containing  the  lesser. 
Selenic  acid  is  dibasic,  and  forms  primary  and  secondary  salts, 
MH  Se  04  and  M2  Se  04 . 

The  properties  of  telluric  acid  are  much  like  those  of  selenic 
acid,  with  the  exception  that  the  anhydride  Te  03  is  known.  The 
latter  compound  is  produced  by  heating  telluric  acid,  which  breaks 
down  into  its  anhydride  and  water,  H2  Te  04  =  Te  O3  +  H2  0.  The 
oxygen  compounds  of  the  sulphur  family  are  exactly  parallel  with 
those  of  the  halogen  family  as  regards  stability.  Thus,  fluorine 
form  no  oxide,  and  the  oxide  of  oxygen  (ozone)  is  quite  unstable ; 
chlorine  forms  a  greater  variety  of  oxides  than  does  any  other 
element  of  the  halogens,  while  sulphur  exhibits  the  same  prop- 
erty in  its  own  family.  The  third  members  of  either  family, 
selenium  and  bromine,  are  the  elements  which,  next  to  the  first 
member  of  each  family,  have  the  least  tendency  to  form  oxides ; 
while  with  iodine  and  tellurium  the  capacity  to  form  stable  oxides 
is  once  more  manifested. 

Both  selenium  and  tellurium  can  form  compounds  with  sulphur ; 
but,  while  those  of  selenium,  excepting  the  simplest  (the  mono- 
sulphide,  Se  S),  are  not  very  definitely  understood,  those  of  tellu- 


162       OXIDES   OF   SULPHUR   FAMILY;   THERMOCHEMISTRY   OF. 


rium  exactly  correspond  to  the  oxygen  compounds  of  that  element, 
so  that  we  have  a  disulphide,  Te  S2  ,  and  a  trisulphide,  Te  S3  ;  and 
this  latter  compound  resembles  the  trioxide,  Te  03  ,  so  far  as  to  be 
the  anhydride  of  an  acid,  for  with  sulphides,  like  those  of  sodium  or 
potassium,  it  will  form  salts  in  which  all  of  the  oxygen  is  replaced 
by  sulphur  :  — 

K2S  +  TeS3  =K2TeS4, 
=  K2Te04. 


The  existence  of  such  a  compound  as  K2  Te  S4  is  another  proof  of 
the  great  resemblance  between  sulphur  and  oxygen. 

The  chlorine  compounds  of  selenium  and  tellurium  are  like 
those  of  sulphur. 

The  following  table  contains  the  thermo-chemical  relations  which 
exist  between  the  oxides  of  the  sulphur  family  and  between  the 
acids,  HX03,  of  the  halogen  family.  Unfortunately,  the  heats  of 
formation  of  the  oxides  of  the  halogens  have  not  been  accurately 
determined,  so  that  in  this  table  we  are  compelled  to  compare  ox- 
ides with  compounds  of  oxygen  and  hydrogen  :  — 


STABILITY. 

S02 
Se02 
Te  O2 

710  K 
572  K 
773  K 

H  Cl  O3 
HBr03 
HI03 

239  K* 
160  K* 
582  K* 

x 

HEATS  OF  FORMATION  OF  THE  CHLOBIUES. 

STABILITY. 

S2C12 
Seo  C12 

143  K 
222  K 

A 

SeCl4 
Te  C14 

462  K 
774  K 

The  heats  of  formation  and  stability  of  the  chlorine  compounds 
increase  with  the  atomic  weights,  and  therefore  increase  with  the 
decreasing  not-metallic  character  of  the  elements.  This  rule  is 
reversed  with  the  hydrogen  compounds,  so  that,  in  this  family,  the 
chlorine  and  the  hydrogen  compounds  are  the  less  easily  decom- 
posed, the  greater  the  chemical  contrast  between  the  elements 
which  form  them. 

*  Dissolved  in  water- 


NITROGEN;    HISTORY.  163 


CHAPTER   XXIII. 

NITROGEN   AND   THE  ATMOSPHERE. 

Nitrogen,  symbol,  1ST;  atomic  weight  14.03;  specific  gravity,  air 
=  1,  is  .97209,  H2  =  2,  is  28 ;  1  c.c.  at  0°  emd  .76  m.  we^fo 
.00125718  0r-ram.  The  Atmosphere,  specific  gravity,  1 ;  specific 
gravity,  H2  =  2,  *s  28.8 ;  1  c.c.  air  at  0°  awd  .76  m.  pressure 
weighs  .00129327 


As  the  atmosphere  is  simply  a  mixture  of  gases,  with  nothing 
in  its  deportment  from  which  general  chemical  comparisons  could 
be  made,  it  seems  advisable  to  treat  of  it  in  a  separate  chapter ;  and 
subsequently  to  introduce  a  discussion  of  the  characteristics  which 
the  elements  of  the  nitrogen  family  have  in  common.  A  descrip- 
tion of  the  element  nitrogen  must,  however,  be  given  before  the 
atmosphere  is  considered. 

An  English  chemist,  Rutherford,  noticed  in  1772  that  a  closed 
space  of  air,  in  which  an  animal  had  breathed  for  a  sufficient  length 
of  time,  was  then  unfit  further  to  support  either  combustion  or  life, 
even  after  care  had  been  taken  to  remove  all  of  the  carbon  dioxide 
by  means  of  lime  water  or  caustic  potash ;  and  from  this  fact  he 
drew  the  conclusion  that  the  residuum  contained  a  peculiar  kind  of 
air.  When,  subsequently,  it  was  discovered  that  elements  which 
form  solid  oxides,  when  burned  in  a  closed  space  of  air,  leave 
nothing  but  this  same  gas,  the  general  consensus  of  opinion  among 
chemists  was  that  this  gas  must  be  phlogisticated  air  j  for  the  burn- 
ing substance  had  given  up  its  phlogiston^  and  that  portion  of  the 
atmosphere  which  remained  must,  as  a  consequence,  have  taken  it 
up  ;  oxygen,  in  which  substances  burned  so  readily,  was  therefore 
most  certainly  pure  dephlogisticated  air.  These  theories  as  regards 
the  nature  of  the  atmosphere  were  unfortunate  for  the  older  school 
of  chemists,  for  it  was  not  difficult  to  prove  that  carbon,  in  burning 
in  oxygen  (or,  as  they  said,  in  giving  up  its  phlogiston),  formed  a 
kind  of  phlogisticated  air  which  was  different  from  the  gas  which 


164     NITROGEN;  OCCURRENCE,  PREPARATION. 

we  now  call  nitrogen,  but  which  was  likewise  unable  to  support 
either  life  or  combustion.  Lastly,  hydrogen,  which  was  supposed 
to  be  pure  phlogiston,  when  burned  in  dephlogisticated  air  forms 
water ;  therefore  that  liquid  should  also  be  phlogisticated  air,  yet 
it  differs  entirely  from  both  of  the  other  forms.  As  a  consequence 
of  these  facts,  chemists  were  brought  face  to  face  with  the  neces- 
sity of  assuming  the  existence  of  three  different  kinds  of  phlogisti- 
cated air.  It  was  left  for  Lavoisier  to  prove  the  fallacy  involved  in 
these  theories  and  to  show  that  the  atmosphere  contained  two  dis- 
tinct gases,  one  of  which  he  called  oxygene,  the  other  azote  (from  a 
and  L'amx6g,  to  sustain  life).  The  name  nitrogene  (from  nitrum, 
saltpetre,  and  the  root  yfv,  to  produce)  was  introduced  by  Chaptal, 
and  from  this  the  English  name,  nitrogen,  has  resulted.  The  word 
azote  is  still  used  by  French  chemists. 

Nitrogen  occurs,  as  such,  in  the  atmosphere,  and  small  quanti- 
ties of  it  are  also  dissolved  in  water ;  combined,  it  is  found  in  the 
form  of  sodium  nitrate  or  Chile  saltpetre,  large  deposits  of  which 
occur  in  the  northern  provinces  of  Chile ;  and  it  also  is  present  as 
ammonium  compounds,  the  latter  being  found  in  small  quantities 
in  the  air  and  in  the  soil ;  nitrogen  is  also  an  essential  constituent 
of  many  organic  substances  of  both  animal  and  vegetable  origin. 
Despite  its  wide  distribution,  nitrogen  forms  scarcely  one  per  cent 
of  the  total  substance  of  the  globe,  even  if  we  include  in  this  its 
gaseous  envelope,  for  very  little,  if  any,  nitrogen  is  found  in  the 
older  geologic  formations. 

The  preparation  of  nitrogen. 

While  we  are  acquainted  with  no  simple  method  by  which  to 
remove  the  nitrogen  from  the  atmosphere  so  as  to  obtain  the 
oxygen  contained  therein,  the  removal  of  the  oxygen  is  a  com- 
paratively simple  matter ;  to  do  this  it  is  only  necessary  to  burn 
some  substance  in  a  closed  volume  of  air.  In  preparing  nitrogen, 
it  is,  of  course,  expedient  to  combust  a  body  which  will  form  a 
solid  oxide,  and  which  will  therefore  leave  no  gaseous  residuum, 
excepting  nitrogen.  Phosphorus  answers  this  purpose  admirably, 
for  when  it  is  burned  it  forms  phosphorus  pentoxide,  a  solid  readily 
soluble  in  water.45  Another  method  for  preparing  nitrogen  is  by 
passing  air  over  copper  foil  heated  in  an  infusible  glass  tube  by 
means  of  a  combustion  furnace ;  the  copper  then  becomes  oxidized, 
copper  oxide,  Cu  0,  being  formed,  while  the  unchanged  nitrogen 
passes  on,  to  be  collected  over  water. 


THE  ATMOSPHERE;  COMPOSITION.  165 

Nitrogen  is  a  colorless,  odorless  gas,  with  a  specific  gravity,  air 
being  one,  of  .97209,  and  hydrogen  being  two,  of  28,  so  that,  as 
the  atomic  weight  of  the  element  is  14,  the  molecule  of  nitrogen 
consists  of  two  atoms,  just  as  is  the  case  with  oxygen,  hydrogen, 
and  chlorine.  Its  critical  temperature  is  —  146°. 3  at  a  pressure 
of  35  atmospheres.  This  means'  that  at  —  146°. 3  nitrogen  becomes 
fluid  if  a  pressure  of  35  atmospheres  is  exerted,  but  that  above  this 
temperature  pressure  cannot  condense  it  to  a  liquid.  The  boiling 
point  of  liquid  nitrogen  is  at  —  194°,  with  a  pressure  of  740  m.  m.  ; 
the  freezing  point  is  at  —  214°.  Nitrogen  is  scarcely  soluble  in 
water,  for  one  volume  of  the  latter  dissolves  only  .015  volumes  of 
the  gas  at  ordinary  temperatures. 

Chemically,  nitrogen  is  a  remarkably  indifferent  substance  when 
it  is  in  the  free  state ;  this  circumstance  is  very  striking  when  we 
consider  that  the  same  element,  when  combined  with  others,  can 
take  part  in  a  great  number  of  chemical  reactions,  and  so  manifest 
the  greatest  chemical  activity.  Nitrogen  will  unite  with  other 
elements  only  under  the  greatest  provocation ;  so,  for  instance, 
nitrogen  and  hydrogen  will  form  ammonia^  under  the  influence  of 
an  electric  current,  and  nitrous  acid,  nitric  acid,  and  ammonium 
nitrate  are  formed  in  small  quantities  when  hydrogen  burns  in  air. 
Owing  to  its  chemical  indifference,  nitrogen  will  neither  burn  nor 
will  it  support  combustion  ;  and,  naturally,  animals  are  asphyxiated 
when  placed  in  the  gas. 

Toward  the  latter  part  of  the  eighteenth  century,  when  chemists 
began  to  inquire  into  the  nature  of  the  atmosphere,  the  determina- 
tion of  the  exact  quantities  of  nitrogen  and  oxygen  going  to  form 
that  gas  became  an  interesting  subject  for  research.  Lord  Caven- 
dish first  made  a  series  of  accurate  investigations  of  the  air  in  and 
about  London,  and  came  to  the  conclusion  that  there  were  about  79 
out  of  every  100  volumes  of  air  left  as  phlogisticated  air  after 
combustion  had  taken  place.  Lavoisier,  of  course,  also  studied  the 
subject,  but  with  very  uncertain  results,  for  at  one  time  he  gave 
the  volume  of  oxygen  as  one-fourth,  at  another  time  as  one-fifth  of 
the  total  atmosphere.  Gay  Lussac  and  Humboldt  made  some  really 
accurate  determinations  of  air  in  the  neighborhood  of  Paris,  and 
they  found  that  there  were  between  20.9  and  21.2  volumes  of  oxy- 
gen in  every  100  volumes  of  air ;  and,  subsequently,  a  great  number 
of  chemists,  notably  Dumas,  Regnault,  and  Bunsen,  made  a  long 


166  THE  ATMOSPHERE ;   CARBON  DIOXIDE  IN. 

series  of  accurate  determinations,  the  results  of  which  showed  that 
the  atmosphere  contained  20.9  volumes  of  oxygen  and  79.1  volumes 
of  nitrogen  in  every  100  volumes,  but  that  these  quantities  were 
subject  to  frequent,  but  very  slight,  variations.  The  following 
figures  will  show  the  extent  of  these  differences  :  — 

Lyons,  Berlin,     20.9  volumes  of  oxygen  in  100  volumes  of  air. 
Algiers,  20.4        "         "       "         "  100        "         "    " 

Calcutta,  20.3         "        "       "         "  100         "         "    " 

Atlantic  Ocean,  21.5        "         "       "        "  100        "         "    " 

These  variations,  which  are  too  great  to  be  within  the  range  of 
experimental  error,  indicate  that  the  atmosphere  is  a  mechanical 
mixture,  and  not  a  chemical  compound.  Another  proof  of  this 
conclusion  can  be  obtained  by  examining  the  difference  in  the  solu- 
bility of  oxygen  and  of  nitrogen  in  water.  A  given  volume  of 
water  will  dissolve  quite  a  little  more  oxygen  than  it  will  nitrogen, 
so  that,  if  we  place  some  water  which  has  been  exposed  to  the  air, 
under  the  bell  of  an  air  pump  and  then  exhaust,  the  bubbles  of 
gas  which  pass  from  the  liquid  will  be  a  mixture  of  oxygen  and 
nitrogen,  containing  more  of  the  former  element  than  does  the  air. 
The  gas  so  formed  can  again  be  dissolved  in  water,  and  the  water 
once  more  exhausted ;  and  if  the  operation  is  repeated  often  enough, 
with  the  same  gas,  the  latter  will  finally  be  almost  pure  oxygen. 
The  oxygen  and  nitrogen  of  the  air  can  therefore  be  separated  by 
simple,  not-chemical,  means.  The  discovery  that  oxygen  and  nitro- 
gen can  be  mixed  to  form  air,  without  a  change  either  in  the  total 
volume  or  in  the  amount  of  heat  contained  in  the  two  gases,  forms 
a  final  argument  against  the  conception  of  the  atmosphere  as  a 
chemical  compound.46 

Oxygen  and  nitrogen  form  the  bulk  of  the  atmosphere,  but 
other  substances  are  always  present  in  minute  quantities.  The 
most  important  of  these  impurities  are  carbon  dioxide,  ammonium 
carbonate,  nitrate  and  nitrite,  and  water  vapor,  as  well  as  solid 
particles  of  dust,  which  are  both  organic  and  inorganic  in  their 
origin.  Carbon  dioxide,  which  is  invariably  found  in  the  air,  is  as 
important  to  living  organisms  as  oxygen  itself.  It  is  of  far  greater 
value  than  is  the  uncombined  nitrogen,  for  the  latest  investigations 
show  that  the  admixture  of  the  latter  gas  does  not  make  the  air 
more  adaptable  for  respiration,  animals  being  able  to  live  as  com- 
fortably in  pure  oxygen  as  they  do  in  air;  on  the  other  hand, 


ATMOSPHERIC   CARBON   DIOXIDE;    REACTIONS.  167 

carbon  dioxide  is  a  plant  food,  and  is  absolutely  necessary  for 
vegetable  life.  Carbon  dioxide  is  being  constantly  added  to  the 
atmosphere  from  burning  fuel,  from  volcanic  craters  and  fissures, 
as  a  product  of  the  breathing  of  animals  and  from  decomposing  or- 
ganic matter,  because  it  is  produced  by  the  combustion  or  decay  of 
carboniferous  substances.  If  no  means  were  provided  for  the 
removal  of  the  atmospheric  carbon  dioxide,  the  increase  in  the 
amount  of  the  latter  would  soon  destroy  all  living  organisms  de- 
pendent upon -respiration.  Fortunately,  plants  growing  in  the  sun- 
light absorb  carbon  dioxide  from  the  air,  using  for  this  absorption 
a  green  coloring  matter  which  is  contained  in  the  leaves,  and  which 
can  eliminate  oxygen  from,  and  add  hydrogen  to,  carbon  dioxide. 
By  this  means  a  substance  which  is  able  to  form  all  of  the  innu- 
merable compounds  of  carbon,  hydrogen,  and  oxygen  occurring  in 
the  vegetable  kingdom,  is  produced.* 

The  carbon  dioxide  of  the  atmosphere  is,  therefore,  continually 
being  removed  while  oxygen  is  being  returned ;  but  were  this  pro- 
cess to  go  on  without  any  compensating  production  of  carbon  diox- 
ide, plant  life,  and  consequently  animal  life,  would  soon  cease. 
The  supply  of  carbon  dioxide  is,  however,  renewed  by  one  means 
or  another,  so  that  the  quantity  in  the  atmosphere  remains  quite 
constant.  In  former  geologic  periods  the  atmosphere  was  undoubt- 
edly much  warmer  and  contained  much  greater  quantities  of  carbon 
dioxide  than  it  does  now.  Hot  rains  were  continually  pouring 

*  According  to  the  theory  proposed  by  Baeyer,  and  which  is  held  by  most 
chemists,  this  substance  is  formic  aldehyde,  CH2  O ;  or 

0  (  =  H2 


. =0  •  (  =  0 

Formic  aldehyde  can  be  considered  as  carbon  dioxide  in  which  one  atom  of 
oxygen  has  been  replaced  by  two  of  hydrogen ;  it  therefore  contains  the  ele- 
ments, hydrogen  and  oxygen,  in  exactly  the  proportions  necessary  to  form 
water,  while  sugars,  starch,  and  cellulose  also  contain  the  elements  in  the 
same  proportions.  Glucose  has  the  formula  C6  H12  O6  =  6  (C  H2  O),  so  that  it 
can  very  possibly  be  formed  by  simple  condensation  of  six  formic  aldehyde 
molecules ;  and,  indeed,  a  substance  very  nearly  identical  with  glucose  has  been 
made  artificially  by  this  means.  Cane  sugar,  starch,  and  cellulose  are  made 
from  C6H12O6  by  separation  of  water  between  the  molecules;  so  that  the 
theory  of  the  reduction  of  carbon  dioxide  by  plants  to  form  formic  aldehyde 
seems  very  reasonable.  The  hydrogen  for  this  reduction  is,  probably,  fur- 
nished by  the  decomposition  of  water,  the  oxygen  of  which  is  eliminated. 


168  ATMOSPHEKIC   CARBON  DIOXIDE  ;   REACTIONS. 

down  and  copiously  watering  the  continents  and  islands  which  had 
been  formed,  so  that  vegetation  on  a  gigantic  scale  nourished 
wherever  the  soil  was  favorable.  Our  coal  beds  were  produced  by 
the  destruction  of  the  flora  of  that  period  ;  by  this  means  enormous 
quantities  of  carbon  dioxide  were  removed  from  the  atmosphere, 
and  the  carbon  stored  for  use  when  the  supply  of  that  gas  should 
no  longer  be  sufficient  to  support  vegetation.  After  the  atmosphere 
had  assumed  the  composition  which  it  possesses  at  present,  animal 
life  flourished.  The  plants  take  up  carbon  dioxide  from  the  atmos- 
phere, form  their  tissues  therefrom,  animals  live  on  them  or  prey 
on  each  other ;  both  plants  and  animals  die,  decay,  and  the  carbon 
dioxide  once  more  finds  its  way  into  the  atmosphere ;  so  that  a  con- 
tinuous metamorphosis,  with  its  energy  given  by  the  light  and  heat 
of  the  sun,  is  in  progress.  Small  quantities  of  carbon  dioxide  are, 
however,  lost  in  the  formation  of  carbonates  of  the  metals,  because, 
in  disintegrating,  the  silicates,  which  form  the  main  body  of  the 
rocks,  liberate  the  bases  which  they  contain,  so  that  these  can  unite 
with  other  acids  (of  which  carbonic  acid  is  one),  and  thus  a  certain 
amount  of  the  latter  substance  is  permanently  removed  from  the 
atmosphere.  This  loss  has  been  supplied  by  the  carbon  dioxide  of 
volcanic  origin  and  by  that  formed  in  burning  the  fuel  which  was 
stored  as  coal  in  a  former  geologic  era ;  so  that,  as  far  as  we  know, 
the  total  amount  of  carbon  dioxide  in  the  atmosphere  is  not  dimin- 
ishing ;  if  it  is  growing  less,  the  rate  of  decrease  is  so  very  slow 
that,  in  the  short  time  during  which  chemists  have  been  able  to  make 
accurate  observations,  no  change  could  be  noted.  The  quantity 
of  carbon  dioxide  in  a  given  volume  of  the  air  varies  slightly ;  but, 
normally,  it  is  about  three  parts  in  ten  thousand,  and  it  seems  that 
the  proportion  of  carbon  dioxide  is  greater  at  night  than  in  the 
daytime,  and  in  summer  than  in  winter.  In  the  higher  regions  of 
the  atmosphere,  where  vegetation  is  impossible,  the  amount  of  the 
gas  may  even  increase  to  eleven  parts  in  ten  thousand,*  while  a  con- 
tinued rainstorm  may  diminish  it  to  two  and  a  quarter.  The 
amount  of  carbon  dioxide  in  crowded  rooms  is  increased  by  the 
breathing  of  the  people  within  the  closed  air  space,  yet  this  does 
not  generally  take  place  to  such  an  extent  that  the  oppressive  feel- 
ing caused  by  such  an  atmosphere  can  be  ascribed  entirely  to  it; 
the  unpleasant  effect  is  due  to  exhalations  of  organic  matter  which 
pass  from  the  lungs.  The  presence  of  carbon  dioxide  in  the  atmos- 

*  Doubtful. 


THE  ATMOSPHERE;   WATER   VAPOR  IN.  169 

phere  can  be  proved  at  any  time  by  exposing  some  clear  lime  water 
to  the  action  of  the  air,  for  a  white  crust  of  the  carbonate  of  cal- 
cium will  be  formed  in  a  short  time.47 

Water  vapor  is  always  present  in  the  atmosphere  in  quantities 
varying  with  the  temperature,  season  of  the  year,  and  locality ;  it  is 
just  as  important  as  carbon  dioxide  to  living  organisms.  The  evap- 
oration of  oceans,  lakes,  and  rivers  furnishes  a  never-ending  supply 
of  water,  the  amount  of  which  is  generally  greater  in  hot  than  in 
cold  weather,  and  greater  by  day  than  by  night. 

The  higher  the  temperature  of  a  gas,  the  greater  will  be  the 
amount  of  water  vapor  which  it  can  take  up,  for  the  quantity  of  the 
latter  which  can  be  contained  in  a  closed  space  (either  a  vacuum  or 
filled  with  gas)  increases  with  the  temperature,  but  is  an  unalterable 
amount  at  any  definite  point.  If  as  much  water  vapor  as  can  possi- 
bly be  present  at  the  existing  temperature  is  contained  in  a  gas,  the 
latter  is  said  to  be  saturated  with  that  vapor,  and  no  further  evapo- 
ration of  water  can  take  place  unless  the  temperature  is  increased ; 
on  the  other  hand,  a  decrease  would  diminish  the  amount  of  vapor 
which  can  be  present,  so  that  a  portion  of  the  moisture  would  be 
precipitated  as  water.  From  these  considerations  it  follows  that,  if 
the  atmosphere  is  nearly  saturated  with  moisture,  any  diminution 
in  the  temperature  will  cause  a  fall  of  rain  or  the  formation  of  dew, 
while  at  the  same  time  no  evaporation  can  take  place  when  such 
a  condition  prevails.  The  amount  of  water  in  the  atmosphere  is 
generally  greatest  near  the  seashore ;  for,  owing  to  changes  in  tem- 
perature, much  of  the  water  will  be  precipitated  before  the  moisture- 
laden  air  can  pass  far  inland.  Drops  of  water  collect  on  a  cold  sur- 
face, because  the  air  in  the  immediate  neighborhood  is  cooled  below 
the  point  at  which  it  is  saturated  with  vapor.  This  point  is  called 
the  dew  point ;  and,  as  the  exact  amount  of  water  vapor  which  can 
be  contained  in  a  given  volume  of  the  atmosphere  at  any  definite 
temperature  is  known,  the  discovery  of  the  dew  point  affords  a 
ready  means  of  ascertaining  the  amount  of  moisture  in  the  air. 
The  ratio  between  the  tension  of  the  water  vapor  which  would  be 
found  were  the  air  fully  saturated  at  the  prevailing  temperature, 
and  that  tension  which  really  exists,  is  called  the  relative  humid- 
ity.* The  quantity  of  water  present  in  the  atmosphere  can  also  be 

*  By  vapor  tension  at  a  given  temperature  is  meant  the  pressure,  in  milli- 
meters of  the  barometer,  which  is  exerted  by  a  vapor  at  that  temperature. 


170  THE   ATMOSPHERE  ;    AMMONIA   IN. 

ascertained  by  passing  a  known  volume  of  air  over  weighed  tubes 
filled  with  calcium  chloride  or  any  other  substance  which  will 
readily  absorb  moisture,  for  then  the  gain  in  weight  will  give  the 
exact  amount  of  moisture  which  is  present. 

The  water  in  the  atmosphere  is  absolutely  essential  to  plant  life. 
The  liquid  falls  upon  the  soil  as  rain,  and  is  then  absorbed  by  the 
radicles;  afterward  it  circulates  through  the  entire  system  of  the 
plant,  taking  part  in  various  physiological  changes,  and  finally 
evaporates  from  the  leaves.  The  amount  of  moisture  which  passes 
from  large  areas  covered  by  vegetation  is  enormous,  so  that  wooded 
districts  cause  an  equitable  distribution  of  rain. 

Another  impurity  present  in  the  atmosphere  is  ammonia;  this 
substance,  however,  is  always  found  combined  with  acids,  as  ammo- 
nium carbonate,  nitrate,  or  nitrite.  These  salts  are  washed  into  the 
soil  by  the  rain,  and  are  then  taken  up  by  plants  to  form  those  por- 
tions of  their  tissues  which,  in  addition  to  carbon,  hydrogen,  and 
oxygen,  also  contain  nitrogen,  so  that  the  ammonium  compounds  in 
the  atmosphere  are  essential  to  plant  life ;  yet  the  amount  of  these 
is  very  small  and  variable  ;  the  greatest  quantity  ever  found  has 
been  47.6  parts  by  weight  in  one  million  of  the  atmosphere.  Small 
quantities  of  other  impurities,  such  as  sulphur  dioxide  and  sul- 
phuretted hydrogen,  may  occur  in  restricted  areas  where  such  gases 
are  being  formed ;  as,  for  instance,  in  districts  where  large  quanti- 
ties of  sulphur-bearing  coal  are  burned.  Ozone  is  also  at  times 
-supposed  to  be  present  in  the  atmosphere. 

The  solid  particles  floating  in  the  air  as  dust  may  be  of  two 
kinds,  inorganic  and  organic.  Sodium  chloride  is  always  present 
in  the  inorganic  particles ;  the  organic  substances  may  be  of  the 
greatest  variety,  and  frequently  contain  micro-organisms  which  can 
inaugurate  disease. 

The  pressure  which  the  atmosphere  is  exerting,  by  reason  of  its 
weight,  is  measured  by  the  barometer.  In  the  seventeenth  century 
some  Florentine  pump-makers,  wishing  to  convey  water  to  a  very 
great  height  by  means  of  a  long  suction  pump,  discovered  to  their 
chagrin  that,  no  matter  how  great  their  exertions,  the  water  would 
not  follow  the  piston  for  more  than  thirty-two  feet,  and  so  Galileo 
Galilei  was  appealed  to  for  an  explanation.  The  cause  assigned  to 
this  phenomenon  by  the  great  man  was,  however,  entirely  a  wrong 
one,  for  he  maintained  that  a  column  of  water  longer  than  thirty- 


THE   BAROMETER.  171 

two  feet  would  be  broken  by  its  own  weight,  just  as  would  a  bar  of 
iron  of  sufficient  length ;  therefore,  water  could  never  be  pumped 
to  any  great  height.  Torricelli,  a  pupil  of  Galileo,  soon  after  (1643) 
found  the  right  explanation.  He  reasoned  that,  when  the  pump 
created  a  vacuum,  the  water  was  pressed  upward  by  the  weight  of 
the  atmosphere;  if  this  were  so,  the  height  of  a  column  of  a  spe- 
cifically heavier  liquid,  such  as  mercury,  which  the  atmosphere  would 
be  capable  of  sustaining,  should  be  proportionally  less.  Accord- 
ingly, Torricelli  rilled  a  glass  tube  (sealed  at  one  end)  with  mer- 
cury, closed  the  open  end  with  his  thumb  and,  inverting  the  tube, 
placed  this  in  a  vessel  filled  with  the  same  metal ;  the  column  of 
mercury  then  sank  until  its  upper  surface  was  between  28  and  29 
inches  above  the  level  of  the  liquid  in  the  trough,  so  that  a  vacuum 
was  produced  in  the  upper  part  of  the  tube.  The  experimenter 
now  concluded  that  the  weight  of  the  column  of  mercury  in  the 
tube  must  be  equal  to  the  weight  of  a  column  of  water  with  an 
equal  cross  section  and  a  height  of  thirty-two  feet,  while  both  col- 
umns exerted  a  pressure  which  was  opposed  by  an  equal  one  pro- 
duced by  the  atmosphere  acting  on  the  surface  of  the  liquid  in  the 
open  vessels  in  which  the  tubes  were  placed.  If  this  conclusion 
was  correct,  the  height  of  a  column  of  mercury  which  could  be  sus- 
tained by  the  atmosphere  would  be  less  on  the  mountain-tops  than 
on  the  low  lands,  and  so  Pascal,  hearing  of  Torricelli's  experiment, 
induced  his  brother-in-law,  Perrier,  to  ascertain,  experimentally,  if 
this  theory  was  really  the  correct  one.  Perrier  took  a  barometer  to 
an  altitude  of  5,000  feet,  and  reported  that  at  that  elevation  the 
mercury  stood  three  inches  lower  than  it  did  in  Paris.  The  whole 
matter  was  now  clear  ;  the  atmosphere  exerted  a  pressure  which 
could  be  measured  by  the  height  of  the  column  of  mercury  it  could 
sustain,  while  the  instrument  constructed  with  this  end  in  view 
subsequently  became  known  as  the  barometer. 

The  distance  from  the  upper  level  of  the  mercury  in  the  barome- 
ter tube  to  that  in  the  vessel  underneath,  is  the  height  of  the  barom- 
eter ;  at  the  level  of  the  sea  the  average  is  760  m.  m.,  and  in  all 
scientific  work  this  is  taken  as  the  standard  for  barometric  meas- 
urements ;  as  the  weight  of  a  column  of  mercury  760  m.  m.  long 
and  with  one  square  inch  cross  section  is  15  pounds,  it  follows  that 
the  pressure  exerted  by  the  atmosphere  is  15  pounds  to  the  square 
inch.  We  do  not  feel  this  enormous  weight  because  the  air  presses 


172  VOLUMES   OF   GASES;   CALCULATIONS. 

equally  in  all  directions,  and  because  the  pressure  from  within  our 
bodies  counterbalances  that  from  without. 

The  volume  occupied  by  any  gas  is  inversely  as  the  pressure 
exerted  on  it.  Double  the  pressure  and  you  halve  the  volume, 
quadruple  the  pressure  and  the  volume  will  be  one-fourth,  and  so 
on.  If  V0  and  V,  be  the  volumes  of  a  gas  at  the  same  temperature 
but  under  different  pressures,  P0  and  P, ,  then :  — 

1.  V0:  V,::P,:P0, 

2.  V0P0=V,P,, 

now  if  P0  is  the  standard  barometric  pressure  of  760  m.  m., 
then :  - 

3.   V0  =  -l^-< 
760' 

or,  the  volume  of  a  gas  at  standard  pressure  is  equal  to  the  volume 
at  the  existing  pressure  multiplied  by  that  pressure  in  millimetres 
and  divided  by  760 ;  but  the  existing  pressure  is  the  pressure  of 
•the  atmosphere  measured  by  the  height  of  the  barometer,  h, 
(P,  =  h)sothat:- 

4        Vn   —  ' 

'    760 

All  gases,  when  not  very  near  the  point  at  which  they  become 
liquid,  expand  ^fa  of  their  volume  for  each  degree  of  temperature,* 
so  that  -a  litre  of  gas  at  0°  will  be  1  -f  ^^  litres  at  +  1°  and 
1  -j_  jfy  at  10°,  and  ten  litres  would  be  10  +  2W  at  +  1°,  conse- 
quently :  — 

«t       6.   V=Vl     «t.         7.    V  =  _L 


where  V0  is  the  volume  of  any  gas  at  0°,  V,  the  volume  of  the 
same  gas  at  any  temperature,  t  is  that  temperature,  and  «  =  ^73. 
Uniting  3  and  7  in  one  equation,  we  have  :  — 

8'  Vo  =  76oJ'+uQ 

•so  that  the  volume  of  any  gas  observed  at  0°  and  760  m.  m.  is 
equal  to  the  volume  multiplied  by  the  height  of  the  barometer  and 
divided  by  760  times  one  plus  ^3-  of  the  temperature.  If  the  gas 

*  More  accurately  .00367. 


THE   ATMOSPHERE;   DEPTH   OF.  173 

to  be  measured  is  in  a  eudiometer  tube,*  partially  filled  with,  mer- 
cury, the  pressure  under  which  it  is,  is  naturally  not  that  of  the 
atmosphere,  but  is  atmospheric  pressure  minus  the  height  of  the 
column  of  mercury  in  the  tube.  If  this  height  in  millimetres  be 
called  w,  then  :  — 

9    y  =    V,(h-w) 

760  (1  +  «  t) 
• 
and,   furthermore,  if  the  gas   is   saturated  with  water-vapor,   the 

tension  of  water-vapor  in  millimetres,  at  the  temperature  of  obser- 
vation, must  also  be  deducted  from  the  barometric  measurement, 

so  that :  —  ,T    .  , . - 

10     V    _  V,  [h  -  (w  +  «fr)] 

760(1 -fat) 

Where  V(  =  volume  observed. 

h  =  observed  barometric  pressure. 

w  =  height  of  mercury  column  above  trough. 

4>  =  tension  of  water  vapor  in  millimetres. 

a   =  .00367. 

t    =  observed  temperature. 

When  the  relations  between  volumes  of  gases  are  considered,  these 
are  supposed  to  be  under  standard  conditions,  at  0°  and  760  in.  m. 
pressure. 

The  height  of  the  barometer  at  any  place  is  constantly  undergo- 
ing variations,  for  the  atmosphere  is  always  subject  to  more  or  less 
serious  local  disturbances,  which  affect  the  pressure  exerted  by  it. 
The  depth  of  the  atmosphere  is  uncertain ;  it  has  been  variously 
given  at  from  thirty  to  two  hundred  miles.  As  the  pressure  is 
greatest  on  the  surface  of  the  earth,  the  air  must  be  densest  at  this 
point,  and  must  diminish  in  density  the  higher  the  altitude.  Prob- 
ably at  an  elevation  of  ten  miles  the  pressure  of  the  atmosphere 
would  be  imperceptible.  The  temperature  of  the  air  becomes  less 
the  greater  the  distance  from  the  earth.  The  specific  weight  of  the 
air  at  760  m.  m,  and  0°,  being  the  most  frequently  used  standard 
of  measurement,  is  generally  taken  as  unity.  Sometimes,  however, 
gases  are  measured  with  hydrogen  as  the  standard,  when  the  specific 
gravity  of  air  becomes  14.38  (hydrogen  =  1),  or  28.76  (hydrogen 
=  2).t  In  order,  therefore,  to  find  the  specific  gravity  of  any  gas, 

*  See  note  20  of  the  Appendix.        t  In  round  numbers,  14.4  and  28.8. 


174  THE   ATMOSPHERE  ;    RELATION   TO   LIFE. 

with  hydrogen  as  unity,  we  must  multiply  the  specific  gravity, 
referred  to  air,  by  14.38. 

All  living  organisms  upon  the  earth  are  dependent  upon  the 
atmosphere.  Its  oxygen,  its  carbon  dioxide,  its  ammonia,  and  its 
water-vapor  are  necessary  to  all  forms  of  life,  for  those  constituents 
which  cannot  be  directly  used  by  animals,  indirectly  find  their  way 
into  their  systems.  The  air  is  inspired  by  animals,  its  oxygen  is 
absorbed,  comes  in  contact  with  every  tissue  in  the  body  and  is  ex- 
haled, charged  with  carbon  dioxide  and  water-vapor,  after  its  oxi- 
dizing action  is  completed.  The  plants  make  use  of  the  carbon 
dioxide  which  finds  its  way  into  the  atmosphere  ;  when  in  the  sun- 
light they  assimilate  it  and  thus  form  the  greater  portion  of  their 
tissues,  but  plants,  as  well  as  animals,  require  oxygen  for  their 
existence;  neither  can  plants  live  without  the  presence  of  com- 
pounds of  nitrogen,  for  many  of  their  most  essential  chemical  con- 
stituents, such  as  the  albumens,  are  composed  chiefly  of  carbon, 
hydrogen,  oxygen,  and  nitrogen.  A  little  of  this  nitrogen  may 
possibly  be  taken  directly  from  that  which  is  contained  in  the 
atmosphere,  but  certainly  the  major  portion  is  furnished  by  com- 
pounds of  nitrogen  found  in  the  soil.  These  compounds  would 
soon  be  entirely  used  up  were  it  not  for  their  constant  renewal  by 
the  addition  of  those  substances  which,  originally  in  the  atmo- 
sphere, are  washed  to  the  ground  by  rains,  and  by  such  nitrogenous 
products  as  are  produced  in  the  soil  by  the  decay  of  organic  sub- 
stances. The  plants  are  thus  able  to  build  their  tissues  from  the 
simplest  inorganic  materials  —  from  carbon  dioxide,  water,  ammo- 
nium nitrates  and  nitrites.  Animals  have  no  such  power ;  they 
destroy,  where  plants  create.  Some  live  upon  plant  substances 
and  assimilate  the  ready-formed,  complicated  organic  compounds  ; 
others  prey  upon  each  other,  so  as  to  get  these  constituents  second- 
hand ;  they  all  die  and,  by  decaying,  once  more  return  to  the  soil 
and  air  those  substances  which  the  plants  had  used  ;  thus  a  cease- 
less rotation  of  the  life-supporting  constituents  of  the  atmosphere 
is  going  on ;  while  the  energy  necessary  to  cause  these  metamor- 
phoses is  furnished  by  the  heat  and  light  of  the  sun. 


ELEMENTS    OF   NITROGEN   FAMILY. 


175 


CHAPTER  XXIV. 

COMPOUNDS   OF  THE   ELEMENTS    OF   THE   NITROGEN   FAMILY. 

The  elements  of  the  nitrogen  family  are  :  — 

1.  Nitrogen,      atomic  weight  14.03 

2.  Phosphorus, 

3.  Arsenic, 

4.  Antimony, 

5.  Bismuth, 

The  same  changes  in  the  physical  characteristics  of  the  ele- 
ments which,  with  increasing  atomic  weight,  are  observed  in  the 
halogen  and  in  the  sulphur  families,  are  repeated  in  that  group  of 
elements  of  which  nitrogen  is  a  representative.  This  will  be  seen 
from  the  following  table  :  — 


O,         N,  colorless  gases. 


Cl,  yellow  gas, 


P,  yellow,  easily  fused  solids. 


Br,  brown  liquid,   Se,        As,  metallic  appearing  solids. 


[,  black  solid,          Te,       Sb,  silver-white  appearing  solids 


-      Bi,  reddish  metallic  solid. 


Specific  Melting 

gravities.  points. 


HYDROGEN  COMPOUNDS. 


General  formula  of  hydrogen  compounds  of 

halogens,  H  X. 

General    formula  of   hydrogen    compounds 

FH 

0    H2 

N    H3 

of  sulphur  group,  H2  X. 

General  formula  of  hydrogen  compounds  of 

C1H 

S    H2 

P    H3 

nitrogen  group,  H3  X. 

Valence  of    the    elements    of    the    halogen 

Br  11 

SeHa 

AsH3 

family  toward  hydrogen  is  1. 

Valence  of  the  elements  of  the  sulphur  fam- 

I H 

TeHa 

SbH3 

ily  toward  hydro'gen  is  2. 
Valence  of  the  elements  of  the  nitrogen  fam- 

ily toward  hydrogen  is  3. 

In  the  nitrogen  family,  as  well  as  in  the  other  two  which  we 
have  studied,  the  elements,  as  the  atomic  weights  increase,  change 

*  An  element,  not  as  yet  discovered,  belongs  in  the  interval  between  anti- 
mony and  bismuth ;  as  a  consequence  the  differences  between  antimony  and 
bismuth  are  much  more  marked  than  are  the  differences  between  any  other 
neighboring  pair  of  elements  we  have  so  far  considered. 


176 


ELEMENTS    OF    NITROGEN    FAMILY. 


into  substances  entirely  metallic  in  appearance,  and  the  alteration 
is  even  more  fundamental  in  its  character  than  it  is  in  the  family 
of  the  halogens  or  in  the  oxygen  family,  for  the  chemical  as  well  as 
the  physical  properties  of  the  last  two  elements  in  this  group  are 
more  metallic  than  they  are  not-metallic.  Once  more  we  come  in 
contact  with  a  family  of  elements  in  which  the  one  having  the 
smallest  atomic  weight  is  a  colorless  gas,  that  with  next  higher  a 
yellow  solid,  easily  fused  and  easily  burned,  and  the  next  a  grayish 
white  solid,  with  the  appearance  of  a  metal  and  the  chemical  be- 
havior of  a  not-metal  (for  in  this  respect  arsenic  is  entirely  like 
tellurium).  Antimony  and  bismuth  are  metals  both  in  their  phys- 
ical and  chemical  properties,  but  antimony  shows  a  not-metallic 
nature  in  some  of  its  compounds,  while  bismuth  is  always  a 
metal.  The  melting  points  of  the  elements  in  this  family  increase 
with  the  atomic  weights,  exactly  as  is  the  case  in  the  sulphur  and 
halogen  families. 

All  of  the  elements  belonging  to  the  nitrogen  group,  with  the 
exception  of  bismuth  (which  is  too  much  of  a  metal  to  do  so),  form 
gaseous  hydrogen  compounds  which,  following  the  general  rule, 
diminish  in  stability  as  the  atomic  weight  of  the  characterizing 
element  increases.  These  compounds  are  formed  by  union  of  three 
hydrogen  atoms  to  one  of  the  nitrogen-like  element  —  so  that  their 
general  formula  is  H3  X,  while  the  valence  of  the  typical  element 
is  three.  By  comparing  all  the  elements  in  the  halogen,  sulphur, 
and  nitrogen  groups,  we  can  see  that,  as  the  atomic  weights  of 
the  families,  as  a  whole,  diminish,  the  valences  toward  hydrogen 


increase. 


F.      0.      N. 
At.  wt.     19        16      14 

F  H.        O  H2        N  H3 

As  we  pass  from  one  element 
to  another  corresponding  one  to 
the  right  of  it,  there  is  a  dimi- 
nution of  not  more   than  five 
units  in  the  atomic  weight  and 
an  increase  of  one  in  the  val- 
ence.   No  elements  exist,  the 
atomic  weights  of  which  lie  be- 
tween any  two  on  the  horizontal 
lines,  so  that  these  twelve  ele- 
ments form  a  portion  of   the 
natural    grouping  obtained  by 
arranging  the  elements  in  the 
order  of  their  atomic  weights. 

Cl.       S.     P. 
At.  wt.    35.5    32      31 

C1H.       SH2        PH3 

Br.      Se.    As. 
At.  wt.    80        79      73 

Br  H.       Se  Ha      As  H3 

I.       Te.  Sb. 

At,  wt.  127      125     120 

IH.         TeH2      SbH3 

Valence  I.           II.           III. 

Atomic  wts.  ^~^^> 

Valence. 

ELEMENTS    OF   NITROGEN   FAMILY.  177 

The  hydrogen  compounds  of  the  elements  of  the  nitrogen  fam- 
ily show  a  chemical  character' which  differs  from  that  found  in  the 
other  two  groups  which  have  been  studied,  and  the  reason  for  this 
difference  is  to  be  found  in  the  greater  number  of  hydrogen  atoms 
which  are  joined  to  one  atom  of  the  typical  element  contained  in 
them.  The  hydrogen  compounds  of  the  halogens  are  acids,  be- 
cause one  atom  of  each  of  those  elements  can  unite  with  but  one 
atom  of  hydrogen;  the  positive  character  of  the  latter,  therefore, 
is  not  sufficient  to  counterbalance  the  negative  properties  of  the 
halogen.  The  elements  of  the  sulphur  family  form  hyrlr.Qgfm.-p*>m- 
pounds  which  are  but  slightly  acid ;  for  the  two  hydrogen  atoms 
contained  in  these  have  twice  the  effect  of  the  one  in  the  chlorine 
group,  while,  lastly,  the  hydrngp.Ti  p.n^pnnnrls  of  the  elements  of 
the  nitrngftTT_jFfljm1y  a.rp.  p.it.hftT  ha.sip.  nr^Tiftiitrfl.l,  for  here  the  three 
hydrogen  atoms  entirely  counterbalance  the  chemical  character  of 
that  of  the  element  with  which  they  are  united.  However,  an 
atom  of  hydrogen  has  a  small  mass,  so  that  its  influence  on  the 
character  of  a  compound  would  become  less  as  the  mass  of  the 
atom  of  the  element  with  which  it  is  united  becomes  greater. 
This  connection  between  mass  and  chemical  character  can  be  seen 
in  the  sulphur  and  nitrogen  families ;  thus,  water  is  nearly  neutral 
in  behavior,  being  both  basic  and  acid  ;  sulphuretted  hydrogen 
(where  the  mass  of  the  sulphur  atom  is  twice  that  of  an  oxygen 
atom)  has  the  power  of  reddening  litmus,  while  its  basic  character 
is  much  less  than  that  of  water ;  lastly,  hydrogen  selenide  and 
telluride  are  also  weakly  acid.  In  the  nitrogen  family,  ammonia 
is  a  pronounced  base,  being  capable  of  uniting  with  almost  all  acid 
substances  to  form  stable  compounds ;  PH3  (phosphine)  yields  salts 
with  but  a  few  acids,  such  as  H  Br  and  HI,  while  As  H3  (arsine) 
and  SbHs  (stibine)  can  form  no  salts ;  so  that  the  increase  in  the 
mass  of  the  not-metallic  elements  has  gradually  counterbalanced 
the  effect  of  the  positive  hydrogen,  even  in  spite  of  the  fact  that 
the  not-metallic  character  of  the  elements  themselves  has  dimin- 
ished. Were  antimony  as  negative  as  nitrogen,  stibfrie  would,  in 
that  event,  undoubtedly  be  an  acid ;  it  being  the  diminishing  not- 
metallic  character  of  the  elements  in  a  family,  as  we  pass  to  those 
members  with  higher  atomic  weights,  which  influences  the  character 
of  the  hydrogen  compounds  which  are  formed.  Should  iodine,  to 
cite  another  example,  be  as  not-metallic  as  chlorine,  we  would 


178      HYDROGEN   COMPOUNDS    OF   NOT-METALS ;    COMPARISON. 

expect  hydroiodic  acid  to  be  a  much  stronger  acid  than  hydrochloric, 
for  the  mass  of  an  atom  of  iodine  is  greater  than  that  of  an  atom 
of  chlorine ;  that  this  is  not  the  case  is  due  to  the  fact  that  the  in- 
creased mass  of  iodine  renders  that  element  less  not-metallic  in  its 
character  than  is  chlorine.  Organic  chemistry  teaches  us  that  if 
we  substitute  a  more  positive  group  of  elements  for  the  hydrogen 
contained  in  ammonia,  the  resulting  compound  becomes  more  like  a 
metal ;  and  that  by  substituting  a  less  positive  group,  it  becomes 
less  like  a  metal  in  its  character;  and  also  that  arsine  and  stibine, 
which  are  neutral,  can  be  made  to  act  like  metals  if  we  only 
replace  their  hydrogen  atoms  by  groups  of  atoms  having  a  more 
positive  chemical  character.* 

Whatever  truth  there  may  be  in  the  above  considerations,  the 
facts  themselves  are  to  be  remembered  ;  so  that :  — 

Compounds  HX  of  the  halogens  are  acid. 

Compounds  H2X  of  the  sulphur  group  are  acid  (and  sometimes 
basic). 

Compounds  H3  X  of  the  nitrogen  group  are  basic  or  neutral. 

In  calling  these  compounds  acid,  we  mean  that  they  react  with 
bases  to  form  salts  or  salt-like  bodies,  in  calling  them  basic,  that 
they  can  unite  with  acids  to  form  similar  substances.  The  follow- 
ing reactions  will  make  this  clear :  — 

HC1+     KOH  =  KC1  +  H90. 
H2S  +  2  KOH  =  K2  S  +  2  H20. 
Acid.         Base. 
HC1+     H3N^=NH4C1. 
H2S  +     H3N  =  NH4SH. 
Acid.         Base. 

Water  can  sometimes  act  like  a  base.  This  fact  is  seen  from  its 
behavior  when  it  is  brought  in  contact  with  anhydrides  and  is  made 
apparent  by  comparing  the  following  two  equations :  — 

S03-+-H20     =     H2S04, 

S03+K20     =     K2S04.t 
Anhydride.     Base.        Salt. 

*  A  "positive"  element  is  one  which  in  a  given  chemical  compound 
behaves  like  a  metal;  a  negative  element  is  one  which  behaves  like  a  not- 
metal  ;  while  positive  and  negative  groups  are  such  as  can  chemically  show  the 
same  contrast  to  each  other  as  do  a  metal  and  a  not-metal. 

t  See  pages  30  and  31. 


ELEMENTS    OF    NITROGEN    FAMILY;    OXIDES. 


179 


One  great  difficulty  in  attempting  a  classification  is  found  in  the 
necessity  of  using  such  arbitrary  terms  as  "  basic "  and  "  acidic," 
which  are  often  differently  applied  by  different  chemists ;  this  diffi- 
culty becomes  less,  for  the  present,  if  we  remember  that  only  the 
general  and  important  chemical  characteristics  shown  by  the  hydro- 
gen compounds  of  the  not-metals  are  at  present  being  considered ; 
so  that  when  the  student  subsequently  discovers,  for  instance,  that 
the  hydrogen  atoms  in  ammonia  can  be  replaced  by  potassium,  he 
need  not,  for  that  reason,  look  upon  ammonia  as  an  acid  —  he  need 
remember  merely  that  this  is  an  individual  case,  which  is  an  ex- 
ception to  the  rule.  (See  page  75.) 

All  of  the  elements  of  the  nitrogen  family  form  oxides  ;  the  gen- 
eral formulae  of  the  most  important  of  these,  together  with  those  of 
the  acids  derived  from  them,  are  given  on  the  following  table ;  the 
oxides  of  the  corresponding  halogen  compounds  are  also  tabulated 
for  purposes  of  comparison  :  — 


HALOGENS. 

VAL- 
ENCE. 

NITROGEN  FAMILY. 

In  the  nitrogen  family, 
the  acids  H  Y  O2  and  H  Y  O3 
can  (excepting  the  nitrogen 
acids)  form  stable  hydrated 
acids  :  — 
HYO2  +  H2O  =  H3YO3 
HY03  +  H2O  =  H3YO4 

OXIDES. 

ACIDS. 

OXIDES. 

ACIDS. 

NOMENCLATURE. 

X20 
X203 

x,p. 

(X.  07) 

HXO 
HX02 
HXO  3 
HX04 

1 
3 
5 

7 

Hypo-ous  acid 
ous  acid 
ic  acid 

Y203 
»Q, 

II  Y02 

H  Y  o3 

In  this  table  X  represents  any  halogen,  Y  any  member  of  the  nitrogen 
family. 

The  trioxides  and  the  pentoxides,  as  well  as  the  acids  derived 
from  them  by  the  addition  of  water,  have  corresponding  formulae 
in  both  families,  and  the  nomenclature  of  the  acids  is  also  parallel, 
thus :  — 

C1203  yields  HC102  (chlorous  acid),  and  N203  yields  HN02, 
(nitrous  acid). 

I205  yields  HI03  (iodic  acid),  and  jST205  yields  HN03  (nitric 
acid). 

There  are  no  very  important  acids  in  the  nitrogen  family  corre- 
sponding to  hypochlorous  and  hypobromous  acids ;  and,  as  the 
highest  valence  of  the  elements  of  this  group  toward  oxygen  is 
five,  there  can  foe  none  corresponding  to  perchloric  acid,  which 
is  derived  from  chlorine  with  a  valence  of  seven.  All  of  the  ele- 
ments of  the  nitrogen  family,  with  the  exception  of  nitrogen  and 


180  ELEMENTS   OF  NITROGEN   FAMILY;   OXIDES. 

bismuth,  have  a  great  tendency  to  form  hydrated  acids  which  are 
much  less  readily  decomposed  than  are  those  of  the  corresponding 
class  in  the  other  two  families.  As  a  consequence,  phosphorous 
acid  exists,  not  as  HP02,  but  as  H3P03  (HP02  +  H2  0),  and 
phosphoric  acid  as  H3  P04  (HP03  -+-  H2  0),  more  often  than  as 
HP03,  so  that  the  acids  most  frequently  met  with  contain  the 
same  number  of  hydrogen  atoms  in  a  formula  weight  as  there  are 
hydrogen  atoms  in  a  molecule  of  ammonia.  The  oxide  Y203  be- 
comes more  basic  as  the  atomic  weight  of  Y  increases,  so  that, 
while  N2  08  and  P2  03  are  the  anhydrides  of  acids,  As2  03  is  both  a 
base  and  an  anhydride,  Sb203  is  more  basic  than  acidic,  while 
Bi2  03  acts  altogether  as  a  base.  This  change  in  the  nature  of  the 
oxides  is  the  natural  result  of  the  change  from  not-metal  to  metal 
which  takes  place  in  this  family,  as  we  pass  from  member  to 
member  in  the  direction  of  increasing  atomic  weights. 

The  pentoxides,  Y2  05 ,  all  are  anhydrides 
of  acids  with  no  basic  properties  ;  *  therefore, 
an  addition  to  the  amount  of  oxygen  present 
in  the  compounds  which  are  basic,  changes 
the  latter  into  more  negative  and,  as  a  conse- 
quence, acidic  bodies.  The  general  rule  is 
that,  where  several  oxides  of  the  same  metal 
exist,  the  character  of  these  becomes  less 
basic  as  the  number  of  oxygen  atoms  in  a 

molecule  increases;  so  that  frequently  the  lowest  oxide  may  be  a 
strong  base,  forming  most  stable  salts  with  acids,  while  the  highest 
may  be  the  anhydride  of  an  acid,  and,  as  a  consequence,  the  oxides 
of  the  same  element  may  not  resemble  each  other  so  much  as  they 
do  the  corresponding  oxides  of  some  other  element. 

Nitrogen,  in  addition  to  the  two  oxides,  N2  03  and  oST2  05 ,  forms 
three  others  with  the  formulae,  N2  0,  NO,  and  N02 ;  these  will  be 
discussed  at  the  proper  time.  The  chlorides  of  the  elements  of 
the  nitrogen  family  have  the  general  formulae  Y  C13  and  Y  C15 ; 
their  stability  increases  as  the  metallic  character  of  the  elements 
becomes  more  pronounced  :  — 
N  C13  (?),  explosive. 

*  Bi2  O5  shows  scarcely  any  of  the  properties  of  an  anhydride ;  the  com- 
pound H  Bi  O3 ,  corresponding  to  HNO3 ,  has  been  isolated,  but  it  forms  no 
salts  and  shows  a  tendency  to  break  down,  giving  off  oxygen. 


BASIC. 


ELEMENTS    OF   NITROGEN   FAMILY;   CHLORIDES.          181 

P  C13 ,  P  C15 ,  decomposed  by  water  to  phosphorous  and  phospho- 
ric acids. 

As  C13 , ,  exists  in  the  presence  of  little  water,  entirely 

decomposed  by  much  water. 

Sb  C18 ,  Sb  C15 ,  partially  decomposed  by  water. 

BiClg, ,  partially  decomposed  by  water. 


182  AMMONIA;    OCCURRENCE,   HISTORY. 


CHAPTER   XXV. 

AMMONIA     AND     THE    OTHER  COMPOUNDS   OF   NITROGEN 
AND  HYDROGEN. 

Ammonia;  formula,  NH3.  Specific  gravity,  air  =  \,  is  .589,  H2 
=2,  is  16.96  ;  molecular  iveight,  17  ;  1  cc.  at  0°  and  .  76  m.  pres- 
sure, weighs  .0007651  gram.  Hydrazin ;  formula,  N2  H2 . 
Azoimid;  formula,  N3H. 

AMMONIA  is  by  far  the  most  important  of  the  three  compounds 
of  nitrogen  and  hydrogen.  It  is  never  found,  as  such,  in  nature, 
but  always  occurs  as  an  ammonium  salt,  in  combination  with  some 
acid  ;  in  the  atmosphere  and  soil  it  is  found  as  ammonium  carbon- 
ate, nitrite,  and  nitrate ;  in  mineral  waters  and  in  volcanic  regions, 
as  the  sulphate.  Ammonium  compounds  occur  in  almost  all  plants, 
in  the  air  exhaled  from  the  lungs,  and  in  the  urine  of  animals. 
Those  ammonium  compounds  which  are  found  in  the  soil  and  in 
clay  are  produced  from  outside  sources. 

Salts  of  ammonium  were  first  introduced  into  Europe  from 
Eastern  countries,  especially  from  Armenia,  so  that  the  name  given 
by  the  Arabian  alchemists  to  the  chloride,  the  salt  which  earliest 
came  into  prominence,  was  sal  armeniacum  ;  but  this  term  was  sub- 
sequently altered  to  sal  ammoniacum,  which  had  been  given  to  so- 
dium chloride  imported  from  the  neighborhood  of  the  temple  of 
Jupiter  Ammon  in  the  Libyan  Desert.  Ammonium  compounds 
were  first  prepared  in  the  East  by  the  distillation  of  camels'  dung, 
and  were  much  prized  as  universal  remedies ;  subsequently  other 
animal  refuse  was  used  in  the  preparation  of  the  carbonate ;  at  one 
time  the  substance  prepared  by  the  dry  distillation  of  harts'  horns 
was  considered  a  most  potent  medicine,  so  that  the  name,  spirits  of 
hartshorn,  is  still  used  to  designate  a  solution  of  ammonia  in 
water.  No  distinction  was  made  between  ammonia  and  the  car- 
bonate of  ammonium ;  both  were  called  volatile  alkali,  until 
Priestley  collected  pure  ammonia  gas  over  mercury.  The  term 
ammoniaque  was  introduced  by  the  French  chemists  at  the  close  of 
the  last  century,  and  from  this  our  term  ammonia  has  its  origin. 


AMMONIA  ;    PREPARATION.  183 

Preparation  of  ammonia  by  reduction  of  the  oxides  of  nitrogen. 

Under  ordinary  circumstances  nitrogen  and  hydrogen  have  no 
tendency  whatever  to  unite ;  only  if  electric  sparks  are  allowed  to 
pass  for  some  time  through  a  mixture  of  the  two  gases  does  union 
take  place.  If,  however,  hydrogen  acts  upon  some  oxide  of  nitro- 
gen under  proper  conditions,  especially  if  the  hydrogen  is  in  the 
so-called  nascent  state,  ammonia  is  formed  with  the  greatest  ease. 
For  instance,  if  a  mixture  of  nitric  oxide  (NO)  and  hydrogen  are 
passed  through  a  tube  heated  to  dull  redness  and  containing  a  few 
pieces  of  spongy  platinum,  ammonia  and  water  are  produced  :  — 

NO  +  5H  =  NII3-hH20. 

If  pieces  of  iron  or  zinc  act  upon  very  dilute  nitric  acid,  am- 
monium nitrate  is  produced ;  the  hydrogen  which  first  results  from 
the  action  of  the  metal  on  the  acid  not  only  takes  away  all  of  the 
oxygen  from  the  compound  of  nitrogen  which  it  attacks,  but  in 
addition,  even  unites  with  that  element  to  form  ammonia.  (See 
nitric  acid.) 

Preparation  of  ammonia  from  organic  substances. 

Decaying  organic  substances,  which  contain  nitrogen,  give  off 
ammonia,  and  this  substance,  uniting  with  the  carbon  dioxide 
formed  at  the  same  time,  produces  ammonium  carbonate.  Owing 
to  its  formation  in  this  way  the  odor  of  ammonia  can  always  be 
detected  in  the  neighborhood  of  heaps  of  manure  or  in  stables. 
The  compounds  of  nitrogen,  which  are  an  essential  constituent  of 
existing  plants  and  animals,  must  have  been  just  as  necessary  in 
past  geologic  eras,  so  that  the  vegetable  remains,  which  are  found 
in  the  form  of  bituminous  coal,  are  rich  in  compounds  of  carbon, 
hydrogen,  oxygen,  and  nitrogen;  the  dry  distillation  of  this  coal 
at  present  produces  nearly  all  of  the  ammonia  in  use.  When 
bituminous  coal  is  heated  without  the  access  of  air,  a  large  num- 
ber of  products  of  great  commercial  importance  are  given  off. 
They  are :  — 

Gaseous  —  Illuminating  gas,  ammonia,  sulphuretted  hydrogen. 

Liquid  —  Water,  benzol,  toluol,  phenol,  etc. 

Solid  —  Naphthalene,  anthracene,  etc. 

The  highest  boiling  products  are  substances  like  asphalt,  while 
the  carbon  remains  behind  in  the  form  of  coke. 

The  ammonia  which  passes  from  the  gas  manufacturing  retorts 


184 


PREPARATION. 


is  separated  from  the  mixture  of  gases  (obtained  by  the  dry  distil- 
lation of  coal)  by  passing  the  gases  upon  surfaces  of  wood  over 
which  a  continuous  stream  of  water  is  trickling ;  *  the  solution  so 
formed  is  a  dark-colored  liquor  from  which  it  is  an  easy  matter 
to  obtain  ammonia  in  a  pure  state,  t  To  do  this,  the  dissolved 
ammonia  can  be  converted  either  into  ammonium  chloride  or  sul- 
phate by  the  addition  of  hydrochloric  or  sulphuric  acid.  The  so- 
lution of  ammonia  in  water,  we  may  assume,  for  purposes  of 
comparison,  contains  the  gas  combined  with  water  as  ammonium 
hydroxide :  -  NH3  +  H2  0  =  KH4  OH. 

When  acids  are  added  to  this  hypothetical  hydroxide,  an  ammonium 
salt  is  produced,  just  as  a  potassium  or  sodium  salt  would  be  formed 
under  similar  circumstances  :  — 

KOH+    HC1  =         KC1        +     H90, 

2KOH  +  H2S04  =      K2S04      +2H20, 

2s»(NH4)OH  +   HC1   =    (NH4)C1    +     H20, 

2  (NH4)  OH  +  H2  S04  =  (NH4)2  S04  +  2  H20, 

(The  group  of  elements  represented  by  the  formula  NH4  very  much 
resembles  potassium.)  In  order  to  liberate  the  ammonia  from  these 
salts,  it  is  only  necessary  to  treat  them  with  some  such  base  as  so- 
dium or  calcium  hydroxide,  for  these  not-volatile  bases  would  expel 
a  volatile  one  from  its  salts,  so  that :  - 

NH4C1     +    NaOH     =    NaCl  +     NH4OH, 

(NH4)2  S04  -f-  2  Na  OH   =  Na,  S04  +  2  NH4  OH, 

2(NH4)C1     +  Ca(OH)2  =    Ca C12  +  2  NH4 OH. 

The  ammonium  hydroxide  which  would  be  formed  breaks  down 
into  ammonia  and  water,  NH4  OH  =  NH3  -f  H2  0,  so  that  pure  am- 
monia gas  can  be  liberated  from  the  chloride  or  sulphate  of  ammo- 
nium. The  gas  is  then  collected  by  passing  it  into  water,  and  the 
solution  so  formed  is  the  liquor  ammonia  of  commerce.  A  similar 
method  can  be  employed  in  the  production  of  small  quantities  of 

*  The  greater  portion  of  the  illuminating  gas  is  insoluble,  or  nearly  in- 
soluble in  water.  The  ammonia,  on  the  other  hand,  is  quite  readily  soluble, 
so  that  it  can  be  separated  from  the  other  constituents  of  the  gas  by  solution 
in  water. 

t  The  solution  contains  the  least  part  of  the  ammonia  in  the  form  of  free 
ammonia;  the  greater  part  is  in  the  form  of  ammonium  salts,  such  as  ammo- 
nium sulphide,  cyanide,  or  sulphocyanide. 


AMMONIA;   PROPERTIES. 


185 


ammonia  for  laboratory  use.48  One  objection  to  the  method  of 
obtaining  pure  ammonia,  which  has  been  detailed,  is  found  in  the 
fact  that  the  sulphate  and  chloride  so  produced  are  very  impure 
and  permeated  by  a  tar-like  substance.  As  a  consequence,  of  late 
years,  the  crude  liquors  are  directly  distilled,  after  the  addition  of 
lime  (the  latter  substance  being  used  to  decompose  the  ammonium 
salts  present  in  the  liquid),  the  tarry  matter  being  removed  from 
the  ammonia  by  filtering  through  pieces  of  paraffine.  By  this 
process  of  distillation  the  pure  ammonia  passes  off  first,  and  is 
collected  in  water. 

Ammonia  is  a  colorless  gas  with  a  penetrating  alkaline  odor. 
At  —  40°  and  under  a  pressure  of  one  atmosphere,  it  condenses  to 
a  colorless  liquid  which  boils  at  —  38°. 5  and  which  changes  into  a 
crystalline  solid  at  —  75°.  The  gas  has  a  specific  gravity  of  .589 
at  0°  and  760  m.m.,  which,  with  hydrogen  as  two,  gives  a  density  of 
16.96,  so  that  the  molecular  weight  of  this  compound  is  17.  As 
a  consequence  it  is  specifically  lighter  than  air,  and  can  be  col- 
lected by  upward  displacement,  exactly  as  is  done  with  hydrogen. 
Ammonia  is  quite  a  stable  compound,  just  as  are  the  correspond- 
ing substances  in  the  oxygen  and  halogen  families,  namely,  water 
and  hydrofluoric  acid ;  it  follows  that  ammonia  does  not  support 
combustion.  It  also  burns  with  difficulty,  even  when  it  is  mixed 
with  oxygen  and  ignited ; 49  the  products  of  combustion  are  water, 
nitrogen,  and  traces  of  the  oxides  of  nitrogen.  Naturally  it  does 
not  support  life.  Ammonia  can  be  completely  decomposed  into 
nitrogen  and  hydrogen  by  passing  an  electric  spark  through  the 
gas  for  some  length  of  time;  and,  when  this  is  done,  two  vol- 
umes of  ammonia  yield  three  of  hydrogen  and  one  of  nitrogen. 


= 

H 

1 
H 

(H 

KVH 

(H 

I 

3 

s^ 

r 

H 

1 
H 

(H 

N^H 
(H 

I 
I 

I 
I 

or  2  NH3  =  N2  +  3 


186 


AMMONIA  ;   COMPOSITION   BY   VOLUME. 


Ammonia,  when  it  is  heated  to  780°  in  a  porcelain  or  iron  tube,  is 
almost  completely  dissociated  into  hydrogen  and  nitrogen. 

Ammonia  can,  under  proper  circumstances,  burn  in  oxygen  to 
form  water  and  nitrogen,  and  it  is  acted  011  by  chlorine  in  a  similar 
way,  producing  hydrochloric  acid  and  nitrogen  :  — 

3  +  3  Cl  =  N  +  3  H  Cl. 


By  means  of  this  reaction  we  can  readily  determine  the  relative 
volumes  of  nitrogen  and  hydrogen  which  go  to  make  ammonia.  If 
we  fill  a  long  glass  tube  with  chlorine,  and  then  add  ammonia 
water,  without  admitting  any  air,  hydrochloric  acid  and  nitrogen 
will  be  formed.  We  learned  that  one  volume  of  chlorine  unites 
with  one  volume  of  hydrogen  to  form  hydrochloric  acid,  so  that  the 
volume  of  chlorine  in  the  vessel  must  have  united  with  an  equal 
volume  of  hydrogen,  which  was  contained  in  the  ammonia;  and, 
therefore,  the  volume  of  hydrogen  which  was  present  in  the  ammo- 
nia is  given  by  the  volume  of  chlorine  in  the  vessel.  Now,  by 
cautiously  admitting  very  dilute  sulphuric  acid,  we  can  absorb  the 
excess  of  ammonia,  while  the  hydrochloric  acid  has  already  formed 
ammonium  chloride  with  the  ammonia  water.  Nothing  but  nitro- 
gen will,  therefore,  remain  in  place  of  the  chlorine  originally  used, 
and  its  volume  will  exactly  fill  one-third  of  the  tube.  It  follows 
that  one  volume  of  nitrogen  and  three  of  hydrogen  unite  to  form 
ammonia.  The  following  diagram  will  make  this  clear  :  — 


in  solution  = 


The  hydrochloric  acid  is  removed  by  dissolving  in  ammonia  water, 
so  that  the  nitrogen  remains,  and  this  is  one  third  of  the  original 
volume  of  chlorine. 

We  have  seen  that  when  ammonia  is  decomposed  by  the  electric 
spark,  two  volumes  of  the  gas  yield  three  of  hydrogen  and  one  of 
nitrogen,  and  that,  conversely,  ammonia  is  formed  from  three  vol- 
umes of  hydrogen  and  one  of  nitrogen.  From  the  first  of  these  two 
discoveries  we  can  conclude,  as  we  do  in  respect  to  oxygen  and 


AMMONIA;    PROPERTIES. 


187 


Jf 


chlorine,  that  nitrogen  contains  two  atoms  in  its  molecule  (see  page 
71)  ;  from  both  we  conclude  that  the  formula  of  ammonia  is  NH3 . 
In  addition,  the  specific  gravity  of  ammonia  shows  us  that  its  mo- 
lecular weight  is  17;  and  quantitative  analysis  shows  that  in  17  parts 
by  weight  of  ammonia  there  are  fourteen  of  nitrogen  and  three  of 
hydrogen.  The  atomic  weight  of  nitrogen  is  therefore  presumably 
14,  and  must  remain  so  unless  some  compound  of  nitrogen  should 
be  discovered,  the  molecular  weight  of  which  is  known,  and  which 
contains  relatively  less  than  fourteen  parts  by  weight  of  nitrogen. 

Ammonia  is  very  soluble  in  water ;  one  volume  of  water  dissolves 
813  times  its  own  volume  of  ammonia  gas.50  The  solution  has  the 
odor  of  ammonia  gas,  is  alkaline  (for  it  changes  red  litmus  to  blue), 
and  is  generally  considered  as  containing  the  ammonia  combined 
with  water  in  the  form  of  ammonium  hydroxide,  NH4  OH.*  Upon 
warming  the  solution,  ammonia  is  expelled,  and  this  fact  is  made 
use  of  in  the  preparation  of  artificial  ice.  An  iron  vessel  (B),  con- 
taining ammonia  water,  is  connected  by  pipes  with  another  one 
constructed  in  the  form  of 
a  double-walled  hollow  cyl- 
inder (A),  all  of  the  con- 
nections being  such  that 
no  gas  can  escape  (Fig. 
10).  Ammonia  water  is 
placed  in  the  first  vessel 
and  is  warmed.  The  es- 
caping gas  passing  into  the 
space  between  the  double 
walls  of  the  second  vessel, 
which  is  cooled  by  means 
of  cold  water,  the  pressure 
of  ammonia  and  the  cold 
combined  condense  the  gas 
to  a  liquid.  The  conditions  are  now  reversed,  the  cylinder  (A)  has 

*  According  to  recent  investigations,  it  seems  probable  that  ammonium 
hydroxide  is  not  formed  when  ammonia  dissolves  in  water,  for  all  other  hydrox- 
ide solutions  conduct  electricity  with  as  great  ease  as  salt  solutions;  but  the 
solution  of  ammonia  in  water  has  only  -^  the  conducting  power  of  a  salt  of  the 
same  molecular  weight.  As  all  of  the  salts  of  ammonium  correspond  to  those 
of  potassium,  it  is,  however,  convenient  to  consider  the  solution  of  ammonia 
in  water  as  ammonium  hydroxide. 


Fig.  10. 


188  AMMONIUM   SALTS. 

its  jacket  of  cold  water  removed ;  while  the  first  vessel  (B)  is  cooled; 
the  condensed  gas  boils  and  is  absorbed  by  the  water  contained  in 
B,  and  when  the  liquid  ammonia  boils,  enough  heat  is  absorbed  to 
freeze  a  can  of  water  placed  within  the  hollow  cylinder. 

Ammonia  gas  can  add  itself  to  acids  to  form  salts;  a  few 
examples  of  such  additions  are  represented  by  the  following 
equations :  — 

NH8  +  H  Cl      =  NH4  Cl         (ammonium  chloride), 
2  NH3  +  H2S04  =  (KH4)2S04  (ammonium  sulphate), 
NH3  +  H  N03  =  NH4  ]tf03      (ammonium  nitrate). 

These  salts  are  termed  ammonium  salts.  The  univalent  group  of 
elements  NH4,  which,  chemically,  resembles  potassium,  is  called 
ammonium ;  it  takes  the  place  of  hydrogen  in  the  acids,  and  the 
salts  so  produced  have  a  character  similar  to  that  of  the  salts 
of  the  true  metals.  Where  a  group  of  elements  acts  in  this  man- 
ner, being  transferred  from  one  compound  to  another  without 
decomposition,  just  as  an  element  would  be,  it  is  termed  a  radicle. 
Ammonium  is  the  unchanging  constituent  of  a  large  number  of 
compounds ;  if  the  grouping  of  elements  represented  by  the  formula 
NH4  is  destroyed,  then  ammonium  salts  lose  their  identity  as  such. 
Ammonium  is  a  radicle  which  can  act  like  a  metal ;  we  are  ac- 
quainted with  other  radicles  which  are  composed  entirely  of  iiot- 
metals,  and  which  act  as  much  like  the  latter  as  ammonium  does 
like  the  former ;  indeed,  it  is  very  difficult  to  say  just  which  com- 
pounds shall  be  called  radicles  and  which  shall  not  be.  It  is,  per- 
haps, best  to  limit  the  term  to  such  groups  of  elements  as  are  very 
frequently  met  with  as  the  unchanging  constituents  of  a  large  num- 
ber of  compounds,  and  which  can  be  transferred  from  one  compound 
to  another  as  a  whole,  and  without  alteration. 

When  ammonia  dissolves  in  water,  its  solution,  for  convenience 
in  chemical  notation,  is  generally  considered  as  containing  the 
hydroxide  of  ammonium:  — 

NH3  +  H2  0  =  NH4  OH. 

With  this  hypothesis  in  view,  the  parallelism  between  ammo- 
nium and  the  metals  becomes  more  apparent,  for :  — 


AMMONIUM   SALTS. 


189 


(NH4)OH  +  HC1 

Ammonium  hydroxide  +  Hydrochloric  acid 

KOH  +  H  Cl 

Potassium  hydroxide     +  Hydrochloric  acid 

2f(NH4)OH          +  H2S04 

Ammonium  hydroxide  4-  Sulphuric  acid        ; 

2  KOH  +  H2  S04 

Potassium  hydroxide     +  Sulphuric  acid       = 


C1  +  H20. 

Ammonium  chloride  +  Water. 

*C1  +  H20. 

Potassium  chloride    +  Water. 

:(KH4)2S04          +  2H20. 

Ammonium  sulphate  -f-  Water. 

K2  S04  +  2  H2  0. 

Potassium  sulphate    +  Water. 


Ammonium  salts,  upon  heating,  decompose  into  ammonia  and 
the  corresponding  acid.*  For  example,  ammonium  chloride,  when 
vaporized,  yields  ammonia  and  hydrochloric  acid :  — 

!STH4  Cl  =  NH3  +  H  Cl. 

The  specific  gravity  of  ammonium  chloride  vapor,  if  hydrogen  = 
2,  provided  no  decomposition  had  taken  place,  should  be  the  same 
as  its  molecular  weight,  or  53.5.  In  reality  it  is  only  one-half  of 
this  number,  or  26.75.  Let  us  suppose  a  volume  of  hydrogen 
weighs  two  grams,  tlien  an  equal  volume  of  ammonium  chloride 
vapor  weighs  26.75  grams ;  but  were  it  vaporized  without  decompo- 
sition then  it  would  weigh  53.5  grams ;  it  follows  that  the  ammo- 
nium chloride  has  decomposed  into  a  molecule  of  ammonia  and  one 
of  hydrochloric  acid.f 


1  vol.  H.     1  vol.  NH4C1  vapor. 


2  vols.  (NH.  +  HC1). 

A 


/  53.5 
UTH3 

grams   \ 
+  HClJ 

26.75  gr. 

26.75  gr. 

If  undecomposed.         If  decomposed. 


*  For  exceptions  to  this  rule  see  next  page. 

t  Recently,  ammonium  chloride  lias  been  vaporized  under  less  than 
atmospheric  pressure,  in  which  condition  no  decomposition  took  place,  as  the 
specific  gravity  of  the  vapor  indicated.  The  molecules  of  the  gas  had  the 
formula  NH4C1.  This  proves  that  nitrogen  is  quinquivalent  in  ammonium 
chloride.  The  vapor  density  of  ammonium  chloride  in  an  atmosphere  of  am- 
monia gives  numbers  much  larger  than  those  which  would  have  been  obtained 
had  complete  dissociation  taken  place.  See  Pullinger  and  Gardner;  Proc. 
Chem.  Soc.;  1891,  2. 


190  AMMONIUM  SALTS;  DECOMPOSITION  OF. 

These  conclusions  are  a  necessary  result  of  Avogadro's  hypothe- 
sis (page  71).  Quite  a  number  of  bodies  dissociate  into  two  simpler 
ones  on  vaporizing.  All  of  these  give  abnormal  specific  gravities 
for  their  vapors,  but  the  explanation  is  always  similar  to  that  which 
has  just  been  given  in  connection  with  ammonium  chloride.  The 
fact  that  the  specific  gravity  of  ammonium  chloride  vapor,  if  hy- 
drogen is  two,  does  not  correspond  to  the  molecular  weight,  does  not 
invalidate  the  method  of  obtaining  the  molecular  weights  of  gases 
by  means  of  their  specific  gravities  ;  it  only  shows  us  that  we  must 
be  careful  to  ascertain  if  the  gas,  the  specific  gravity  of  which  we 
are  about  to  determine,  is  identical  in  chemical  constitution  with 
the  liquid  or  solid  from  which  it  is  produced. 

The  atoms  of  nitrogen  are  trivalent  in  ammonia,  but  the  element 
is  unsaturated  (see  page  108),  although  it  is  not  capable  of  taking 
up  any  more  positive  atoms,  such  as  hydrogen,  unless  these  elements 
are  joined  to  a  negative  element  or  group  of  elements.  The  pres- 
ence of  the  large  number  of  hydrogen  atoms  which  are  contained 
in  a  molecule  of  ammonia  has  rendered  that  compound  positive ;  it 
therefore  has  no  tendency  to  further  unite  with  positive  substances  ; 
but  when  a  negative  compound  such  as  hydrochloric  (or  any 
other)  acid  is  added  to  ammonia,  the  latter  separates  the  former 
into  two  parts ;  namely,  into  hydrogen  and  the  not-metallic  ele- 
ment or  group  of  elements  with  which  hydrogen  was  united;  these 
two  parts  add  separately,  so  that  nitrogen  becomes  quinquivalent :  — 

C-K  f  -H 

in  r  _H                   v       _H  in  (  — H                      v       _H 

W   ]    -H  +  HC1=WJ   -H  JI   .     -H  +  HNO.=|n   --H 

1    (  — H                           — H  '    (  —  H                              — H 

L—  ci  L— N03. 

We  have  seen  that  water  is  similarly  decomposed  when  it  adds  it- 
self to  oxides  in  order  to  form  hydroxides  (see  page  117). 

All  ammonium  salts  are  decomposed  by  heat,  with  the  results 
which  are  detailed  under  1  and  2 :  — 

1.  The  ammonium  salts  may  be  entirely  disintegrated,  as  is  the 
case  with  ammonium  nitrate  ;  for,  when  that  substance  is  heated, 
neither  ammonia  nor  nitric  acid  is  produced. 

2.  Ammonia  and  the  acid  are  produced. 

a.    The  acid  may  be  volatile,  then  nothing  remains. 


HYDROXYLAMIN.  191 

b.  The  acid  may  be  not-volatile,  then  it  remains. 

c.  The  acid  may  be  decomposed  by  heat,  then  its  decomposition 

product  remains. 

The  radicle  NH4,  ammonium,  has  never  been  isolated,  but  its 
metal-like  nature  is  shown  by  the  fact  that  it  forms  an  amalgam 
with  mercury.  If  an  ammonium  salt,  like  ammonium  chloride,  is 
decomposed  by  sodium  in  the  presence  of  mercury,  the  ammonium 
liberated  will  form  an  amalgam  with  the  latter.* 

NH4  Cl  +  Na  .=  NH4  +  Ka  CL_ 

The  mercury  expands  and  becomes  of  the  consistency  of  soft 
butter.  Ammonium  amalgam  gradually  decomposes  in  the  air,  am- 
monia and  hydrogen  are  given  off,  and  the  mercury  shrinks  to  its 
former  size.51 

It  has  been  mentioned  (page  183)  that,  when  diluted  nitric  acid 
is  reduced  by  means  of  finely  divided  zinc  or  iron,  ammonium  ni- 
trate is  produced.  Under  other  circumstances  (for  example  when 
the  reducing  metal  is  tin)  reduction  does  not  go  so  far  and  a  par- 
tially oxidized  ammonia,  termed  hydroxylamin,  is  formed.  The 
same  result  can  be  accomplished  by  reducing  nitric  oxide  with  tin 
and  hydrochloric  acid. 

Hydroxylamin  is  ammonia  in  which  one  hydrogen  atom  has 
been  replaced  by  hydroxyl,  as  the  following  structural  formulae 
will  indicate :  — 


H  (OH 

Ammonia.  Hydroxylamin. 

Hydroxylamin  is  basic  in  character,  its  salts  resemble  ammonium 
salts,  and  they  are  probably  similarly  constructed  :  — 

< 

*  Amalgams  are  solutions  of  metals  in  mercury.  They  sometimes  have 
definite  crystalline  forms  or  definite  quantities  of  metal  and  mercury.  The 
mercury  may  in  these  compounds  possibly  play  a  part  similar  to  that  of  water 
of  crystallization. 

t  The  structural  formula  of  hydroxylamin  is  not,  as  yet,  absolutely  cer- 
tain ;  this  doubt  is  due  to  the  fact  that,  heretofore,  it  has  been  difficult  to 
isolate  any  quantity  of  the  pure  base.  This  has,  however,  of  late,  been  suc- 
cessfully accomplished,  so  that  we  may  soon  expect  to  have  all  doubt  removed. 


192  HYDRAZIN. 


OH 
H 
-Cl 

Hydroxylarnin  hydrochloride. 

Free  hydroxylamin  is  a  solid,  crystalline  substance  which  melts  at 
32°-35°  and  which  boils  at  56°-57°  at  22  m.m.  pressure.  When 
heated  at  ordinary  pressures  it  explodes  violently  at  130°.  It  can- 
not be  preserved  for  any  length  of  time  without  decomposing.* 
On  the  other  hand,  the  salts  of  hydroxylamin  are  perfectly  stable. 

Nitrogen  forms  two  other  compounds  with  hydrogen  ;  namely, 
hydrazin,  N2H4,  and  azoimid,  N3H;  both  of  these  substances  were 
recently  discovered  by  Curtius.  Hydrazin  can  be  considered  as 
analogous  to  hydrogen  dioxide,  for  the  latter  is  water  in  which  one 
atom  of  hydrogen  is  replaced  by  hydroxyl  :  — 

H-O-H  and  HO-OH, 

while  the  former  is  ammonia  in  which  one  atom  of  hydrogen  has 
been  replaced  by  the  univalent  group  —  NH2, 

H  —  N  =  H2,  and  H2=N  —  N-=H2; 

the  group  of  atoms  to  which  hydrogen  rs  attached  is,  in  the  one 
case,  —  0  —  0  —  ,  and  in  the  other  =  N  —  N=  ;  as  oxygen  is  bi- 
valent, each  oxygen  atom  in  the  above  group  will  be  capable  of 
uniting  with  only  one  hydrogen  atom  ;  while,  as  nitrogen  is  trivalent, 
each  nitrogen  atom  is  capable  of  further  union  with  two  hydrogen 
atoms.  (See  pages  107,  108.) 

Hydrazin  has,  in  all  probability,  never  been  isolated  in  a  pure 
state  ;  in  reactions  where  it  would  be  expected,  not  it,  but  a  mix- 
ture of  nitrogen  and  ammonia  is  formed.  Hydrazin  is  basic  in  its 
character,  so  that  it,  like  ammonia,  unites  with  acids  to  form  salts. 
These  salts  have  the  same  composition  as  those  of  ammonium,  with 
the  exception  that  one  atom  of  hydrogen  is  replaced  by  the  univalent 
radicle  (the  amido  group)  —  NH2  f  :  — 

*  J.  W.  Briihl;  Berichte  d.  Deutsch.  Chem.  Gesell.;  26,  2508. 

t  Salts  of  hydrazin  in  which  both  —  NH2  groups  were  supposed  to  be  united 
with  acids  (C1H3N  —  NH3C1)  were  formerly  described.  Such  salts  are  now 
known  not  to  exist;  only  one  of  the  —  NH2  groups  can  unite  with  acids. 
Curtius;  Berichte  d.  Deutsch.  Chem.  Gesell.;  26,  409. 


AZOIMID. 


193 


fH 
H 
H 
H 

LCI 

Ammonium  chloride. 


N 


H 
H 
H 


Id 

Hydrazoniurn  chloride 


Hydrazin  is  soluble  in  water  ;  its  solutions  have  an  alkaline  reaction. 
Azoimid,  or  hydrogen  nitride,  N3H,  is  a  colorless  gas  with  a 
peculiar,  very  penetrating  odor  ;  it  is  quite  poisonous,  and  its  solu- 
tion in  water  has  an  extremely  irritating  effect  upon  the  skin.  A 
solution  of  azoimid  can  be  prepared  by  oxidizing  a  solution  of  hy- 
drazin  in  water  by  means  of  nitrous  acid.*  The  most  interesting 
fact  in  regard  to  this  compound  is  that  it  is  a  strong  acid,  greatly 
resembling  hydrochloric  or  hydrobromic  acid,  but  a  short  consid- 
eration will  show  us  the  reason  for  this  chemical  behavior,  for  in 
azoimid  the  mass  of  the  three  nitrogen  atoms  entirely  overbalances 
that  of  the  one  hydrogen  atom,  and  consequently  the  compound,  as 
a  whole,  is  negative  ;  we  would  therefore  expect  azoimid  to  be  acid 
in  its  nature.  f  Azoimid  produces  dense  white  fumes  when  brought 
in  contact  with  ammonia,  just  as  hydrochloric  acid  does  ;  the  sub- 
stance formed  in  the  one  case  is  ammonium  nitride,  just  as  in  the 
other  it  is*  ammonium  chloride  :  — 


=  NH4C1. 

The  solution  of  azoimid  attacks  copper,  aluminium,  zinc,  and  other 
metals,  forming  the  nitrides  and  liberating  hydrogen;  it  dissolves 
oxides  and  hydroxides  of  metals.  The  nitrides  formed,  in  all  cases, 
resemble  the  corresponding  chlorides.  Hydrogen  nitride  differs 
markedly  from  hydrogen  chloride  by  being  very  unstable,  for  even 

*  Curtius;  Berichte  d.  Deutsch.  Chem.  Gesell.;  26,  1263. 

t  As  we  increase  the  relative  number  of  hydrogen  atoms  as  compared  to 
each  nitrogen  atom  in  the  molecules  of  the  three  compounds  of  nitrogen  and 
hydrogen,  we  pass  from  an  acid  to  bodies  with  an  entirely  opposite  chemical 
character.  This  change  reminds  us  of  the  transition  from  not-metals  to  met- 
als which  we  encounter  in  the  natural  families  formed  by  the  elements  which 
we  have  studied  —  with  the  difference  that  with  the  increase  of  the  mass  of 
the  not-metallic  element  in  these  compounds,  the  negative  properties  increase, 
while  in  a  natural  family  they  diminish. 


194  AZOIMID. 

a  shock  or  a  slight  increase  in  temperature  will  cause  it  to  explode 
with  terrific  force.  It  follows  from  this  that  the  grouping  of  three 
nitrogen  atoms  in  this  molecule  takes  place  only  under  great  ten- 
sion, so  that  the  molecule  is  subjected  to  a  constant  strain,  just  as 
a  wound-up  watch-spring  is.  The  structural  formula  assigned  to 
hydrogen  nitride  by  its  discoverer  is  :  — 

N. 

II  )N-H. 

x/ 

The  ammonium  salt  of  azoimide  (NH4N3),  and  the  similar  com- 
pound of  hydrazine  (N5H5),  may  be  mentioned  as  being  two  other 
compounds  composed  exclusively  of  nitrogen  and  hydrogen. 


NITROGEN  ;    COMPOUNDS    WITH    OXYGEN.  195 


CHAPTEK   XXVI. 

THE   COMPOUNDS   OF   NITROGEN   WITH   OXYGEN,   AND  WITH 
OXYGEN   AND    HYDROGEN. 

Nitrous  oxide;  formula,  N2  0 ;  specific  gravity,  air  =  1,  is  1.527, 
H2  =  2,  is  43.98  ;  molecular  weight  is  44.  1  c.c.  of  the  gas  weighs 
.0019835  gram  at  0°  and  .76  m.  pressure.  Nitric  oxide  ;  for- 
mula, N  0  ;  specific  gravity,  air  =  1,  is  1.0384,  H2  =  2,  is  29.9  ; 
molecular  weight  is  30.  1  c.c.  of  the  gas  weighs  .0013488  gram. 
Nitrogen  peroxide ;  formula,  N02 ;  specific  gravity,  air  =  1,  is 
1.58,  H2  =  2,  is  45.5  (at  130°  ),  molecular  weight  is  46  (at  140°  ). 

NITROGEN  forms  the    following   compounds   with   oxygen  and 
hydrogen : — 

1.  No  0,     Nitrous  oxide,  HNO,    hyponitrous  acid, 

2.  NO,      Nitric  oxide,  — -, 

3.  N2  03 ,  Nitrogen  trioxide,  HN02 ,  nitrous  acid, 

4.  NO* ,    Nitrogen  dioxide,  , 

5.  N2  05 ,  Nitrogen  pentoxide,  HN03 ,  nitric  acid. 

Of  these  compounds,  N2  0,  N2  03 ,  and  N2  05  are  similar  to  those  en- 
countered in  the  study  of  the  halogens,  for  there  we  have  C120, 
C12  03 ,  and  L  05 ,  so  that,  provided  we  consider  oxygen  as  being 
uniformly  bivalent,  the  valence  of  nitrogen  in  these  oxides  is  one, 
three,  and  five ;  the  acids  derived  from  these  three  oxides,  HNO, 
hyponitrous  acid,  HN02 ,  nitrous  acid,  HN03 ,  nitric  acid,  also  have 
formulae  like  those  of  the  halogen  acids  ;  but  no  per-nitric  acid  exists, 
so  that  the  highest  valence  displayed  by  the  nitrogen  family  is  five. 
Nitric  oxide  does  not  act  like  the  anhydride  of  an  acid ;  it  is  but 
little  soluble  in  water,  and  is  not  attacked  by  bases ;  neither  has 
nitrous  oxide  the  characteristics  of  an  anhydride,  it  has  no  tendency 
to  form  hyponitrous  acid  with  water  or  hyponitrites  with  alkalies ; 
but,  on  the  other  hand,  it  is  produced  when  a  solution  of  hyponi- 
trous acid  is  warmed,  so  that  it  must  be  looked  upon  as  the  anhy- 
dride of  that  acid.  We  will  begin  the  discussion  of  the  oxides  of 


196  NITROUS  OXIDE;  PREPARATION. 

nitrogen  with  nitrous  oxide,  following  with  nitric  oxide,  nitrogen 
trioxide,  nitrogen  peroxide,  and  nitrogen  pentoxide  in  the  order 
named. 

Nitrous  oxide  never  occurs  as  such ;  it  is  solely  a  product  of  the 
laboratory.  The  gas  was  discovered  by  Priestley  in  1776,  and  was 
first  called  dephlogisticated  nitric  gas ;  its  composition  was  not  ex- 
plained until  some  time  after  its  discovery,  when  Davy  proved  it  to 
be  an  oxide  of  nitrogen.  It  is  best  prepared  by  heating  ammonium 
nitrate,  when  water  and  nitrous  oxide  are  formed  as  follows : 52  — 

NH4  N03  =  N2  0  +  2  H2  0,* 

but  the  gas  can  also  be  produced  by  the  reduction  of  nitric  oxide  by 
means  of  finely  divided  metals,  such  as  zinc,  iron,  or  lead. 

Nitrous  oxide  is  a  colorless  gas,  with  a  very  slight  odor  and 
sweBti^h~tg,ste.  Its  specific  gravity,  air  =  1,  is  1.527,  which,  with 
hydrogen  as  two,  would  give  43.98,  so  that  the  molecular  weight 
is,  in  round  numbers,  44.  In  this  weight,  analysis  shows  that  there 
are  twenty-eight  parts  by  weight  of  nitrogen  and  sixteen  of  oxygen, 
so  that  nitrous  oxide  contains,  in  its  molecule,  one  atom  of  oxygen 
and  two  of  nitrogen  ;  for,  by  means  of  the  study  of  water  and  other 
compounds  of  oxygen,  we  have  concluded  that  the  atomic  weight  of 
oxygen  is  sixteen,  provided  that  of  hydrogen  is  1.008 ;  and  from 
our  study  of  the  composition  of  ammonia  and  other  nitrogen  com- 
pounds, it  follows  that  the  atomic  weight  of  nitrogen  is  14.  Nitrous 
oxide  has,  therefore,  a  structure  similar  to  that  of  water,  as  will 
be  seen  by  comparing  the  formulae :  — 

N  — 0  — N  and  H  —  0  —  H. 

Nitrogen  is,  therefore,  univalent  in  nitrous  oxide,  just  as  hydrogen 
is  in  water.  From  a  further  study  of  the  composition  of  nitrous 
oxide,  we  see  that  two  volumes  of  nitrogen  will  unite  with  one  of 
oxygen  to  form  two  volumes  of  nitrous  oxide,  just  as  two  volumes 
of  hydrogen  unite  with  one  of  oxygen  to  form  two  of  water ;  the 
conclusions  regarding  the  composition  of  water,  at  which  we  arrived 
on  page  73,  are  consequently  equally  applicable  to  nitrous  oxide. 

Nitrous  oxide  is,  to  a  considerable  extent,  soluble  in  water; 
one  volume  of  water  will  absorb  about  1.3  volumes  of  the  gas  at  0° 

*  A  similar  reaction  takes  place  when  ammonium  nitrite  is  heated ;  only 
then  nitrogen  and  not  nitrous  oxide  is  formed :  — 

NH4  NO2  -  2  N  +  2  H.2  O. 


NITROUS   OXIDE  ;   PROPERTIES.  197 

Nitrous  oxide  will  support  combustion  almost  as  readily  as  oxygen ; 
a  glowing  pine  chip  will  take  fire  in  the  gas,  and  phosphorus, 
as  well  as  sulphur,  which  have  been  ignited  in  the  air,  will  con- 
tinue to  burn  brilliantly  in  nitrous  oxide.  The  great  tendency 
to  give  off  oxygen  which  is  displayed  by  nitrous  oxide,  is  readily 
understood  when  we  consider  that  it  is  an  endothermic  compound, 
in  the  formation  of  which  work  which  is  equivalent  to  180  K  must 
be  done;  the  gas,  therefore,  possesses  more  energy  than  its  con- 
stituents, and  will  break  down  at  the  first  opportunity.  That, 
however,  considerable  impulse  is  required  to  inaugurate  this  de- 
composition is  shown  by  the  fact  that  feebly  burning  sulphur  is 
extinguished  in  the  gas,  while  that  which  is  combusting  with 
some  energy  will  continue  to  burn  in  nitrous  oxide  with  almost 
the  same  brilliancy  as  if  it  were  placed  in  oxygen.  When  a  sub- 
stance which,  like  phosphorus,  forms  a  solid  oxide,  burns  in  nitrous 
oxide,  there  is  no  change  in  volume,  for  the  molecule  of  N2  O  sim- 
ply loses  oxygen,  while  a  molecule  of  K  is  left  in  its  place ;  one 
thousand  molecules  of  nitrous  oxide  would  therefore  yield  the  same 
number  of  molecules  of  nitrogen,  or  x  molecules  of  nitrous  oxide 
would  yield  x  molecules  of  nitrogen ;  it  therefore  follows  that  the 
volume  of  nitrogen  which  is  formed  has  the  same  number  of  parti- 
cles as  the  volume  of  gas  from  which  it  is  produced,  provided  the 
gases  have  the  same  temperature  and  are  under  the  same  pressure, 
but,  when  tivo  gases,  under  the  same  temperature  and  pressure,  con- 
tain equal  numbers  of  molecules,  they  have  equal  volumes. 

Nitrous  oxide  is  quite  readily  condensed  to  a  liquid ;  at  0°  it 
becomes  fluid  under  a  pressure  of  thirty  atmospheres ;  its  boiling 
point,  under  atmospheric  pressure,  is  —  88°,  it  becomes   solid  at 
-  100°. 

Although  nitrous  oxide  can  give  up  its  oxygen  so  readily  to 
burning  substances,  it  cannot  do  the  same  thing  in  order  to  support 
respiration.  If  the  gas  is  inhaled,  the  first  effect  is  loss  of  con- 
sciousness, accompanied  by  a  rumbling  in  the  ears,  while  the  person 
undergoing  treatment  experiences  an  involuntary  tendency  to  laugh ; 
as  a  consequence  of  this  effect  Davy  named  this  substance  J.aughinj 
gas.  Small  animals  are  very  rapidly  killed  by  nitrous  oxide.  TEe 
^retfects"of  the  inhalation  of  the  gas  disappear  soon  after  pure  air  is 
taken  into  the  lungs  —  and  as  a  consequence  it  is  extensively  used 
as  an  anaesthetic  in  place  of  chloroform  or  ether.  The  nitrous 


198  NITRIC    OXIDE  ;   PREPARATION. 

oxide  used  for  this  purpose  is  condensed  and  transported  in  iron 
bottles. 

Nitric  oxide  is  the  oxide  of  nitrogen  which  contains  the  next 
greater  quantity  of  oxygen.  It  results  from  the  action  of  many 
metals,  or,  indeed,  of  other  oxidizable  substances  on  diluted  nitric 
acid ;  it  does  not  occur  in  a  free  state,  for  the  oxygen  of  the  atmos- 
phere converts  it  into  a  mixture  of  the  two  higher  oxides,  N2  03 
and  N02 .  The  most  convenient  method  of  preparing  the  gas  is  by 
the  action  of  nitric  acid  on  copper.53  When  copper  is  treated  with 
concentrated  nitric  acid  (spec.  grav.  1.4)  nothing  but  nitrogen  diox- 
ide (  N02 )  and  nitrogen  trioxide  (  N2  03 )  are  produced  as  reduction 
products  of  the  acid.*  These  gases  are  formed  in  the  proportion 
of  about  90  per  cent  of  the  former  and  10  per  cent  of  the  latter. 
As  the  nitric  acid  is  diluted,  the  relative  quantity  of  dioxide  dimin- 
ishes and  that  of  trioxide  increases,  when  the  specific  gravity  of 
the  acid  has  reached  1.25,  nitric  oxide  begins  to  appear,  and,  when, 
more  water  is  added,  nitric  oxide  (mixed  with  a  very  little  nitrous 
oxide)  is  the  sole  product  of  the  reaction.  These  changes  are  very 
readily  understood  if  the  following  facts  are  remembered :  — 

1.  Nitric  oxide  is  oxidized  to  nitrogen  dioxide  and  to  nitrogen 
trioxide  by  means  of  concentrated  nitric  acid.f 

2.  Nitrogen  dioxide,  on  the  addition  of  water,  changes  to  nitric 
acid  and  nitric  oxide  :  — 

3  N02  +  H,  0  =  2HN03  +  NO. 

It  is  very  evident,  therefore,  that  no  nitric  oxide  can  be  formed 
from  concentrated  nitric  acid  and  metals,  so  that  the  first  reaction 
between  the  acid  and  copper  consists  in  a  simple  oxidation  of  the 
metal,  the  nitric  acid  giving  up  the  least  possible  quantity  of 
oxygen,  as  will  be  seen  from  the  equations :  — 

1.  2  HN03  =  2  N02  +  H2  0  +  04 

2.  Cu  +  0  =  CuO. 
Combining  1  and  2  we  have :  — 

3.  2  HN03  +  Cu  =  CuO  +  2  N02  +  H20. 

*  Freer  and  Higley;  Amer.  Chem.  Journ. ;  15,  71. 

t  Veley;  Proc.  Royal  Soc.;  52,  27. 

\  When  nitric  acid  oxidizes  in  such  a  way  that  nitrogen  dioxide  is  the 
reduction  product,  then  two  formula  weights  of  nitric  acid  are  capable  of  fur- 
nishing one  oxygen  atom.  This  fact  must  be  kept  in  mind  in  constructing 
equations. 


NITRIC    OXIDE;   PREPARATION.  199 

One  further  change  must  also  take  place,  because  copper  oxide  is  a 
base,  and,  therefore,  in  the  presence  of  nitric  acid,  must  form  copper 
nitrate  and  water :  — 

CuO  +  2  HN03  =  Cu(  N03  )2  +  H,  0. 

The  complete  reaction  between  copper  and  concentrated  nitric  acid 
is,  therefore,  as  follows  :  — 

Cu  +  4HN03  =  Cu(N03)2  +  2  H20  +  2NO2. 
Now,  if  a  considerable  quantity  of  water  is  present,  then  every 
three  molecules  of  nitrogen  dioxide  change  two  of  nitric  acid  and 
one  of  nitric  oxide ;  and  so  the  following  reactions  take  place  :  — 

a.  (3  Cu  +  12  HN03  =  3  Cu(N03)2  +  6  H20  +  6  N02) 

b.  (6  N02  +  2  H2  0  =  4  HN03  +  2  NO) 

/>!  V   •       •  37  1  C«*   S^   (/ 

Combining  a  and  b  we  have :  — 

3  Cu+12  HN03  +2  H2  0  =  3  Cu  (  N03  )2  +  4  HN03+  2  NO  +  6  H2  0. 

What  is  true  of  copper  is  true  of  the  other  metals  as  well,  with 
the  difference  that  some  of  the  latter  (lead  for  example)  yield 
nitric  oxide,  mixed  with  considerable  quantities  of  nitrous  oxide, 
when  they  act  upon  diluted  nitric  acid. 

As  a  result  of  these  experiences  it  is  expedient  to  prepare  nitric 
oxide  by  the  action  of  copper  on  dilute  nitric  acid. 

Reducing  agents,  which  are  not  of  a  metallic  nature,  such  as 
sulphur  dioxide,  will  also  produce  nitric  oxide  from  nitric  acid. 
For  instance,  dilute  nitric  acid  oxidizes,  dry  sulphur  dioxide  to  sul- 
phuric acid,  while  at  the  same  time  nitric  oxide  is  produced. 
When  nitric  acid  acts  in  this  way,  it  is  convenient  to  ignore  the 
intermediary  production  of  nitrogen  dioxide  and  its  subsequent 
decomposition  to  nitric  acid  and  nitric  oxide,  and  to  consider  the 
oxidizing  action  as  being  always  produced  as  follows : — 

a.  2  HN03  =  H2  0  +  2  NO  +  3  0. 

Two  formula-weights  of  nitric  acid  have,  therefore,  three  atoms 
of  oxygen  at  their  disposal,  and  so  they  can  oxidize  three  molecules 
of  sulphur  dioxide  to  sulphuric  acid  as  follows :  — 

b.  3  S02  +  3  H2  0  +  3  0  =  3  H2  S04. 
Combining  a  and  b,  we  have  :  — 

c.  2  HN03  +  3  S02  +  3  H2  0  =  3  H2  S04  +  2  NO  +  H2  0. 
In  all  other  cases  where  nitric  oxide  is  produced  from  nitric  acid, 


200  NITRIC    OXIDE  ;    PROPERTIES. 

the  reaction  can  be  regarded  as  taking  place  in  a  manner  similar 
to  the  above.* 

Nitric  oxide  is  a  colorless  gas,  which  instantly  turns  dark  brown 
on  exposure  to  the  atmosphere,  nitrogen  trioxide  and  nitrogen  di- 
oxide being  formed ;  as  a  consequence  there  can  be  no  experiment 
showing  whether  the  gas  is  tasteless  and  od.orless  as  well  as  color- 
less. The  two  gases  produced  by  contact  of  nitric  oxide  with  the 
air  are  poisonous.  The  specific  gravity  of  nitric  oxide  is  1.038 
when  air  is  the  standard ;  this  corresponds  to  a  density  of  29.9, 
H  =  2,  or  to  a  molecular  weight  of  30.  It  follows,  as  in  this  molec- 
ular weight  there  are  fourteen  parts  by  weight  of  nitrogen  and 
sixteen  of  oxygen,  that  the  formula  of  nitric  oxide  is  NO.  The 
gas  is  composed  of  equal  volumes  of  nitrogen  and  oxygen,  just  as 
hydrochloric  acid  is  of  hydrogen  and  chlorine.  If  we  wish  to  con- 
sider oxygen  as  bivalent,  then  nitrogen  must  also  be  bivalent  in 
this  compound.  The  specific  gravity  of  nitric  oxide  does  not  change 
even  at  a  temperature  as  low  as  —  70°.  At  —153°. 6  and  at  atmos- 
pheric pressure,  nitric  oxide  changes  to  a  colorless  liquid  which 
solidifies  at  —  167°,  forming  a  snow-like  mass.  The  gas  is  very 
stable ;  it  can  be  heated  to  1200°  without  alteration ;  at  white  heat 
it  is  completely  broken  down  into  nitrogen  and  oxygen.  One  hun- 
dred volumes  of  water  dissolve  about  five  volumes  of  nitric  oxide 
at  ordinary  temperatures. 

Nitric  oxide  does  not  allow  substances  to  burn  in  it  as  readily 
as  does  nitrous  oxide.  For  instance,  phosphorus  does  not  take  fire 
in  the  gas  unless  it  is  heated  to  a  point  considerably  above  its 
melting  temperature ;  in  the  latter  event,  it  will  unite  with  the 
oxygen  of  nitric  oxide  with  the  greatest  energy.  On  the  other 
hand,  sulphur,  a  burning  candle,  or  burning  hydrogen  is  extin- 
guished by  nitric  oxide.  A  mixture  of  carbon  disulphide  and  nitric 
oxide  burns  with  an  exceedingly  brilliant  flame.  Metals  like  zinc 
or  iron,  which  are  easily  oxidized,  will,  if  moist,  readily  remove  a 
part  of  the  oxygen  from  nitric  oxide  and  in  that  way  produce  nitrous 
oxide :  — 

*  All  of  the  methods  which  have  been  given  produce  nitric  oxide  mixed 
with  some  nitrous  oxide.  In  order  to  produce  pure  nitric  oxide,  a  solution  of 
nitrogen  trioxide  in  sulphuric  acid  is  treated  with  mercury.  The  mercury 
will  then  reduce  the  nitrogen  trioxide  to  pure  nitric  oxide.  (See  Emich; 
Monatshefte  f iir  Chemie ;  13,  74. ) 


NITROGEN   TRIOXIDE  ;   NITROGEN  DIOXIDE.  201 

2  NO  +  Fe  =  N2  0  +  Fe  0. 

Priestley  first  prepared  the  latter  gas  by  this  method.*54 

The  existence  of  gaseous  nitrous  anhydride,  N2  03  ,  is  doubtful, 
there  being  strong  reason  to  suppose  that,  in  all  cases  where  chem- 
ists have  endeavored  to  obtain  pure  nitrous  anhydride,  they  have 
only  succeeded  in  producing  a  mixture  of  nitric  oxide  (NO)  and 
nitrogen  dioxide  (  N02  )  ;  such  a  mixture,  obviously,  would  contain 
the  same  proportion  of  nitrogen  and  oxygen  by  weight  as  nitrogen 
trioxide  :  — 

NO  +  N02  =  N2  03 

The  brown  gas  which  results  when  nitric  oxide  is  mixed  with 
an  excess  of  oxygen,  is  nearly  pure  nitrogen  dioxide  (N02);  but  that 
which  is  formed  by  mixing  nitric  oxide  with  ah  amount  of  oxygen 
not  sufficient  to  produce  the  peroxide  undoubtedly  consists  of  a 
mixture  of  nitrogen  dioxide  and  nitrogen  trioxide.  Both  of  these 
gases  are  easily  condensed  to  the  liquid  form,  in  which  state  nitro- 
gen trioxide  certainly  can  exist.  Fluid  nitrogen  trioxide  is  an  in- 
digo-colored liquid  which  is  condensed  at  —  10°  and  which  boils 
below  0°,  giving  off  dark  brown  vapors  which  change,  in  part  at 
least,  into  nitric  oxide  and  nitrogen  peroxide. 

Nitrogen  dioxide  is  produced  when  nitric  oxide  is  exposed  to 
the  atmosphere  :  — 


or  when  the  nitrates  of  certain  metals  are  heated,  lead  nitrate  being 
the  most  convenient  for  this  purpose  :  55  — 

Pb  (N03)2  =  Pb  0  +  2  N02  +  O.f 

*  Nitric  oxide  disolves  in  a  solution  of  ferrous  sulphate,  giving  a  dark  brown 
color.  It  has  been  said  that,  by  heating  this  solution,  pure  nitric  oxide  can 
be  driven  off.  This  is  not  the  case,  however,  as  ferrous  sulphate  reduces  a. 
portion  of  the  nitric  oxide  to  nitrous  oxide,  while  it  is  itself,  in  part,  oxidized 
to  ferric  sulphate.  (Emich;  Monatshefte  fiir  Chemie;  13,  73.) 

t  The  nitrate  of  lead  first,  undoubtedly,  breaks  down  into  lead  oxide  and 
nitric  anhydride:  — 

Pb  (NO8)2  =  Pb  O  +  No  O5  , 
just  as  nitric  acid  would  break  down  into  water  and  nitric  anhydride:  — 


but  the  latter,  at  the  temperature  of  the  reaction,  forms  oxygen  and  nitrogen 
dioxide  :  — 

N2  05  =  2  N02  +  O. 


202    NITROGEN  DIOXIDE  ;  PROPERTIES.      NITROGEN  PENTOXIDE. 

The  gas  has  a  dark  brown  color  which  deepens  as  the  tempera- 
ture is  increased ;  it  has  a  corroding  action,  giving  a  saff ron  coloring 
to  the  skin,  and  other  nitrogen-bearing  organic  compounds.  A 
moderate  cold  condenses  the  gas  to  a  yellow  liquid,  which  becomes 
lighter  in  color  the  lower  the  temperature,  and  which  solidifies  at 
_  9°  to  —  15°  and  boils  at  about  22°.  Nitrogen  dioxide,  when  it  is 
at  a  temperature  just  above  the  boiling  point  of  the  liquid,  has  a 
vapor  density  which  indicates  that  its  molecule  has  the  formula 
N2  04 ;  but  these  molecules,  as  the  heat  is  increased,  begin  to  break 
down  into  those  having  the  composition  N02,  so  that  the  specific 
gravity  of  this  substance  diminishes  until  140°  is  reached,  when 
the  dissociation  of  N2  04  into  N02  is  complete ;  at  600°  the  gas  has 
become  entirely  colorless  and  has  decomposed  into  nitric  oxide  and 
.oxygen.*  Nitrogen  dioxide  is  a  powerful  oxidizer;  carbon  and 
strongly  heated  phosphorus  burn  in  it,  and  the  presence  of  this  gas 
dissolved  in  fuming  nitric  acid  probably  gives  rise  to  the  powerful 
oxidizing  action  of  the  latter  substance.  Nitrogen  dioxide  is 
changed  into  nitric  acid  and  nitric  oxide  when  it  is  dissolved- in 
water  (see  page  198),  so  that  the  same  tendency  to  form  the  acids 
with  the  greatest  possible  amount  of  oxygen,  which  we  observed 
existing  in  the  halogen  and  sulphur  families,  is  once  more  en- 
countered in  the  case  of  the  compounds  under  discussion. 

Nitrogen  pentoxide  is  the  anhydride  of  nitric  acid  and  is  best 
prepared  by  removing  the  wafer  t'rdm~~concentrated  nitric  acid  by 
means  of  phosphoric  anhydride  : 56  — 
02 

r\    •"•'_":  NOo 

DIHJ      =         >0  +  H,0 
iOHi  N02 

o2 

2HN03=     N205     +H20. 

The  compound  is  a  crystalline  solid  which  melts  at  30°  and  boils  at 
45°.5,  while  it  is  at  the  same  time  partially  decomposed ;  it  cannot 

*  Nitric  oxide,  at  a  temperature  above  red  heat,  seems  to  decompose  into 
nitrogen  peroxide  and  nitrogen,  and  the  resulting  nitrogen  peroxide  is  only 
broken  down  completely  at  a  temperature  produced  by  a  white  hot  platinum 
wire.  The  final  result  is  complete  decomposition  into  nitrogen  and  oxygen. 
It  follows  that  the  above  temperature  (600°)  is  open  to  question.  (See  Emich; 
Monatshefte  fur  Chemie;  13,  79.) 


NITRIC    ACID;    HISTORY.  203 

be  kept  for  any  length  of  time  because  of  its  tendency  to  break 
down  into  nitrogen  dioxide  and  oxygen :  — 

N205  =  2N02  +  0. 

Dangerous  explosions  may  be  the  result  of  this  decomposition,  if  the 
pentoxide  has  been  kept  in  a  sealed  tube.  Nitric  anhydride  forms 
nitric  acid  when  it  is  added  to  water :  — 

N205  +  H20=2HN03. 

Nitric  acid  has  been  known  ever  since  the  time  of  the  Arabian 
alchemists.  The  first  authentic  account  of  its  preparation  is  given 
by  Geber,  who,  in  the  ninth  century,  made  it  by  distilling  a  mixture 
of  saltpetre  (potassium  nitrate),  blue  vitriol  (copper  sulphate),  and 
alum  (aluminium  and  potassium  sulphate).  The  first  samples  of 
nitric  acid  were  undoubtedly  an  impure  article.  The  name  given 
to  it  was  aqua  dissolutiva  or  aqua  fortis,  while  the  term  aqua  regia, 
was  used  to  designate  a  mixture  of  nitric  and  hydrochloric  acids. 
Nitric  acid,  or  aqua  fortis,  the  alchemists  discovered,  had  the  power 
of  dissolving  all  known  metals  with  the  exception  of  gold,  while 
aqua  regia  would  attack  even  this,  so-called,  noblest  of  all  metals 
—  almost  nothing  could  withstand  its  corrosive  action ;  surely, 
thought  they,  this  liquid  must  be  closely  allied  to  the  "  alcahest," 
the  universal  solvent  which  they  were  seeking.  At  the  beginning 
of  the  eighteenth  century,  nitric  acid  was  extensively  made  by  the 
action  of  sulphuric  acid  on  nitre  (saltpetre)  ;  to  this  method  of 
preparation  it  owes  its  present  name,  which  is  derived  from  spiritus 
nitri.  Lavoisier  first  proved  that  nitric  acid  contained  oxygen,  and 
its  definite  composition  was  ascertained  during  the  present  century. 

Nitric  acid  can  be  produced  by  the  direct  union  of  nitrogen, 
oxygen,  and  water.  Such  a  synthesis  takes  place  when  electric 
sparks  are  passed  through  moist  air.67  In  all  probability  nitrogen 
peroxide,  N02,  is  at  first  generated,*  however,  the  latter,  when  in 

*  This  method  for  the  preparation  of  nitrogen  dioxide  reminds  us  of  the 
similar  one  used  in  forming  ozone  (see  page  48).  Nitrogen  and  oxygen  are 
both  the  first  members  of  their  respective  families,  there  is  but  little  difference 
between  their  atomic  weights,  and  hence  they  should  show  points  of  resem- 
blance, as  indeed  they  do,  for  they  are  both  colorless  gases.  Ozone  can  be 
considered  as  the  oxide  of  oxygen,  OO2;  it  then  corresponds  to  the  oxides  of 
sulphur,  selenium,  and  tellurium,  SO2 ,  Se  O.2 ,  Te  O2 .  In  the  manner  of  its 
formation  and  in  its  formula  it  is  analogous  to  nitrogen  peroxide,  NO2;  fur- 


204  NITRATES  ;    FORMATION   OF. 

contact  with  water,  breaks  down  into  nitric  oxide  and  nitric  acid. 
(See  page  198.)  Oxides  of  nitrogen  are  also  produced  during  com- 
bustion in  the  air ;  *  these  oxides,  in  contact  with  moisture,  are 
further  converted  into  nitric  acid.  As  ammonia  is  generally  pres- 
ent in  the  atmosphere,  this  substance,  uniting  with  the  nitric  acid, 
produces  ammonium  nitrate,  so  that  this  salt  occurs  in  the  air. 
Nitric  and  nitrous  acids  are,  however,  much  more  readily  formed  by 
the  oxidation  of  ammonia  or  of  the  oxides  of  nitrogen  than  by  the 
direct  union  of  the  elements.  When  organic  substances  (which  con- 
tain nitrogen)  decay,  the  nitrogen  passes  off  as  ammonia,  and  this 
substance,  with  the  acids  present  in  the  air  and  with  the  carbon 
dioxide  formed  at  the  same  time,  produces  ammonium  carbonate, 
nitrate,  and  nitrite  (page  170  ).  When  bases  are  present  in  the  soil, 
an  oxidation  of  the  nitrogen  takes  place  so  that  nitrates  are  produced 
instead  of  ammonium  salts.  Calcium  nitrate  is,  as  a  consequence, 
frequently  found  on  the  walls  of  stables  and  cellars,  while  in  the 
neighborhood  of  East  Indian  villages,  where  the  surface  soil  con- 
tains potash,  potassium  nitrate  is  extensively  met  with;  the  col- 
lecting of  this  substance  forms  the  exclusive  occupation  of  a 
number  of  natives.  Large  deposits  of  sodium  nitrate  occur  in 
the  province  of  Tarapaca.in  the  northern  part  of  Chile,  this  sub- 
stance is  known  as  Chile  saltpetre  or  nitre  ;  its  presence  is  probably 
due  to  the  decay  of  marine  vegetation  which  flourished  on  what  is 
now  terra  firma,  during  the  period  when  a  portion  of  the  South 
American  coast  was  submerged.  This  supposition  is  sustained  by 
the  fact  that  sodium  chloride  and  salts  containing  bromine  and 
iodine  are  found  mixed  with  the  nitre. 

Nitric  acid  is  best  prepared  for  laboratory  use  by  the  addition 
of  sulphuric  acid  to  a  nitrate,  a  method  which  we  have  frequently 
employed  in  the  isolation  of  other  acids  (see  page  152).  The  reac- 
tion may  be  represented  by  the  following  equation :  — 

Na  N03  +  H2  SO,  =  Na  HS04  +  HN03 , 
or,  if  comparatively  little  sulphuric  acid  is  used  :  —          — ' 

thermore,  being  an  endothermic  compound,  it  has  a  great  tendency  to  give 
up  one  atom  of  oxygen,  OO.2  =  OO  +  O,  just  as  nitrogen  peroxide  does, 
N02  =  NO  +  O. 

*  This  formation  takes  place  in  greatest  quantity  when  hydrogen  is 
burned  in  air. 


NITKIC   ACID;   PREPARATION,    PltOPER/TIES.  205 


2  Na  N03  +  H2  S04  =  Na^  S04  +  2  HN03 , 

for,  when  the  quantity  of  salt  to  be  decomposed  is  relatively  great, 
as  compared  with  the  amount  of  acid  used,  then  the  secondary  and 
not  the  primary  sulphate  results.  (See  page  153.)58 

Nitric  acid  is  a  colorless  liquid  which  has  probably  never  been 
prepared  entirely  free  from  water.  It  boils  at  86°,  and  becomes 
solid  at  —  47°  ;  if  it  contains  water  enough  to  have  a  specific  weight 
of  1.3  it  congeals  at  —  19° ;  the  purest  acid  known  has  a  specific 
gravity  of  1.55 ;  *  it  fumes  in  the  air  and  turns  yellow  when  ex- 
posed to  the  sunlight,  because  it  breaks  down  into  nitrogen  peroxide, 
water,  and  oxygen.  The  same  change  takes  place  when  nitric  acid 
is  distilled,  for  the  distillate  from  a  colorless,  pure  acid  is  colored 
because  of  decomposition.  This  behavior  reminds  us  forcibly  of 
the  chlorine  acids.  At  temperatures  just  above  the  boiling  point 
of  nitric  acid,  the  specific  gravity  of  the  vapor  shows  that  but  little 
decomposition  has  taken  place,  for,  while  the  molecular  weight  of 
HN03  would  be  63,  the  specific  gravity  of  the  vapor,  H  =  2,  is  59.3  ; 
the  vapor  density  of  the  acid  diminishes  as  the  temperature  is 
increased,  so  that  at  250°  it  is  36,  therefore,  at  that  temperature, 
the  following  change  has  taken  place :  — 

4  HIST03  =  2  H2  0  +  4  N02  +  02  .f 

Considerable  heat  is  evolved  when  nitric  acid  is  dissolved  in  water, 
so  that  the  dilute  acid  possesses  less  chemical  energy,  and  is  there- 
fore more  stable,  than  the  concentrated  one.  It  is  doubtful  if  defi- 
nite hydrated  acids,  such  as  are  encountered  with  sulphuric  acid,  are 
derived  from  nitric  acid  ;  certainly  the  heat  of  solution  of  the  latter  $ 
is  much  less  than  that  of  the  former.  (See  pages  150,  151.) 

*  The  pupil  must  remember  that  the  specific  gravities  of  liquids  and  solids 
are  taken  with  water  as  unity. 

t  If  a  volume  of  hydrogen  weighs  two  grams,  then  the  same  volume  of 
nitric  acid  weighs  63  grams,  4  volumes  of  nitric  acid  would  therefore  weigh 
252  grams.  These  decompose  into  two  volumes  of  water  vapor  weighing  36 
grams,  4  of  nitrogen  peroxide  weighing  184  grams,  and  1  of  oxygen  weighing 
32  grams.  The  4  volumes  of  nitric  acid  therefore  yield  7  volumes  of  the  de- 
composition products;  these  7  volumes  weigh,  2 52  grams,  1  volume  equal  to 
that  of  2  grams  of  hydrogen  therefore  weighs  36  grams;  in  other  words,  pro- 
vided this  decomposition  takes  place,  the  specific  gravity  of  the  mixed  gases 
must  be  36,  if  H  =  2.  (See  pages  72  and  189.) 

J  71  K  as  compared  with  178  K.  Some  investigators  (Wislicenus,  Ber- 
thellot)  have  maintained  that  definite  hydrated  nitric  acids  exist  in  solution. 


206  NITRIC    ACID  ;    OXIDIZING    ACTION. 

i 
Nitric  acid  has  a  great  tendency  to  give  up  its  oxygen  when  it 

is  brought  in  contact  with  reducing  substances.  Examples  of  this 
oxidizing  effect  we  have  seen  in  the  preparation  of  sulphuric  from 
sulphurous  acid  (page  147),  and  in  the  formation  of  nitric  oxide 
from  copper  and  nitric  acid  (page  199).  Nitric  acid  will  attack 
many  organic  substances,  oxidizing  them,  while  at  the  same  time 
the  acid  itself  is  reduced;  when  the  substance  attacked  is  like 
starch  or  sugar  and  the  acid  is  tolerably  concentrated,  then  the 
oxides,  N203  and  N02,  are  the  main  reduction  products.  The 
organic  substance  is  often  completely  destroyed,  yet  in  quite  a 
number  of  cases  the  body  attacked  is  so  changed  that  the  nitro- 
group,  —  N02  (see  nitrosyl  sulphuric  acid,  page  147,  foot-note),  is 
substituted  for  hydrogen ;  such  an  action  is  produced  when,  under 
certain  circumstances,  nitric  acid  reacts  with  glycerine,  forming 
nitro-glycerine.  Concentrated  nitric  acid  violently  attacks  the  skin 
and  mucous  membrane  ;  that  portion  with  which  it  has  come  in  con- 
tact turns  yellow,  blisters,  and  finally  forms  an  ulcer ;  if  the  acid  is 
somewhat  dilute,  the  yellow  color  will  appear  without  the  blistering. 
Nitric  acid  also  attacks  silk  in  the  same  way,  turning  it  yellow, 
and,  if  the  acid  is  concentrated,  destroying  it;  vegetable  dyes  are 
destroyed  by  it,  so  that  cloth  upon  which  nitric  acid  has  accidentally 
been  dropped  cannot  be  restored  to  its  original  color  by  neutraliza- 
tion with  ammonia  water.69 

As  we  have  seen,  many  metals  dissolve  in  nitric  acid  to  form 
the  corresponding  nitrates,  a  reduction  product  being  produced  at 
the  same  time.  These  reactions  can  practically  be  classed  under 
two  heads. 

a.  Those  in  which  ammonia  is  produced,  the  ammonia  at  once 
uniting  with  nitric  acid  to  form  ammonium  nitrate ;  this  change 
takes  place  when  dilute  nitric  acid  is  added  to  zinc,  tin,  or  to  some 
other  metals;  the  reaction  can  be  represented  in  two  stages,  as 
follows :  — 

1.  8  HN03  +  4  Zn  =  4  Zn  (N03)2  +  8  H  (formation  of  hydro- 
gen). 

2.  HN03-f  8  H=3  H20  +  NH3;    NH3  +  HN03  =  NH4 NO$ 
(reduction  of  nitric  acid  and  formation  of  ammonium  nitrate). 

Uniting  1  and  2  we  have  :  — 

3.  4  Zn  + 10  HN03  =4  Zn  (  N03  )2  +  NH4  N03  +  3  H2  0  (com- 
plete reaction). 


NITRIC    ACID  ;    REDUCTION   OF.  207 

b.  Those  in  which  nitric  oxide  is  formed  by  the  action  of  dilute 
nitric  acid :  — 

3  Cu  +  S  HN03  =  3  Cu  (  N03  )2  +  2  NO  +  4  H2  0.*     (  Page  199.) 

These  two  classes  of  reactions,  however,  only  represent  what 
most  frequently  takes  place ;  it  is  known,  for  instance,  that  when 
zinc  acts  011  a  mixture  of  nitric  and  sulphuric  acids,  a  partially 
oxidized  ammonia  known  as  hydroxylamine,  NH2  OH  (see  page  191), 
results  ;  and  when  copper  is  dissolved  in  diluted  nitric  acid,  nitrous 
oxide,  and  even  nitrogen,  may  be  given  off,  the  production  of 
nitrous  oxide  increasing  with  the  amount  of  copper  nitrate  present. 
In  attempting  to  construct  equations  for  such  reactions  we  there- 
fore generally  represent  only  the  principal  changes  which  take 
place. 

The  reduction  of  nitric  acid  by  metals  is  generally  attributed  to 
the  action  of  nascent  hydrogen,  and,  certainly,  in  some  cases  this 
theory  is  well  founded.  We  cannot  enter  into  the  subject  very  deeply 
in  this  text-book,  f  but  the  following  facts  may  not  be  out  of  place. 
When  a  piece  of  magnesium  is  dissolved  in  dilute  nitric  acid,  hy- 
drogen is  at  first  given  off ;  the  production  of  hydrogen,  however, 
soon  ceases,  while  the  oxides  of  nitrogen  make  their  appearance  ; 
it  also  seems  very  probable  that  the  hydrogen  which  has  been 
occluded  by  palladium  (page  34)  passes  through  nitric  acid,  un- 
changed, until  that  hydrogen  which  is  supposed  to  be  chemically  com- 
bined with  the  palladium  begins  to  be  liberated ;  then  the  evolution 
of  hydrogen  stops,  while  the  lower  oxides  of  nitrogen  make  their 
appearance.  Apparently,  then,  hydrogen  which  is  just  in  the  act 
of  being  liberated  from  its  compounds  has  a  greater  chemical  activ- 
ity than  has  ordinary  hydrogen,  so  that,  whether  we  regard  this 
hydrogen  as  acting  by  reason  of  its  existence  as  individual  atoms 
or  not,  there  is  reason  to  suppose  that  in  some  cases  we  should  con- 
sider the  reduction  of  nitric  acid  by  metals  which  are  dissolving 
therein  as  being  caused  by  hydrogen.  The  equations  given  above 
are  intended  to  illustrate  this  conclusion.  In  the  reactions  which 
have  been  described  as  taking  place  between  copper  and  nitric  acid, 

*  The  reaction  here  given  represents  the  simplest  form  of  the  combined 
equations  a  and  b  on  page  199. 

t  For  a  more  complete  account  of  nascent  reactions  the  pupil  may  refer  to 
M.  M.  Pattison  Muir,  Principles  of  Chemistry.  See  also  page  51. 


208  NITRATES  ;    NITRITES  ;    HYPONITRITES. 

however,  the  evidence  all  seems  to  point  toward  a  direct  oxidation 
of  the  metal  by  the  acid  (see  page  198). 

Nitric  acid  is  a  monobasic  acid  ;  it  has  in  its  formula  weight  but 
one  hydrogen  atom  which  can  be  replaced  by  metals.  The  nitrates 
are  all  decomposed  by  heat,  the  change  taking  place  in  one  of  three 
ways :  — 

a.  The  nitrate  is  entirely  decomposed,  as  is  ammonium  nitrate 
(page  196). 

/?.  The  nitrate  breaks  down  into  the  oxide  of  the  metal,  oxygen 
and  nitrogen  peroxide  :  Pb  (  N03  )2  =  Pb  0  -}-  2  N02  -f  0  (see  page 
201).  If  the  oxide  of  the  metal  is  decomposed  by  heat,  of  course 
nothing  but  the  metal  remains. 

y.  The  nitrate  gives  off  oxygen,  leaving  the  nitrite.  This 
decomposition  is  confined  to  the  nitrates  of  very  pronounced  metals, 
such  as  potassium  or  sodium. 

In  these  decompositions  nitrates  differ  from  chlorates,  for  when 
the  latter  are  heated,  the  perchlorates  are  quite  often  produced ;  the 
reason  is  obvious,  nitrogen  cannot  take  up  more  oxygen  than  is 
necessary  to  form  the  oxide  N2  05 ,  which  is  the  anhydride  of  nitric 
acid  (  page  197),  so  that  no  pernitrates  can  be  formed. 

The  existence  of  nitrous  acid  is  doubtful,  although  the  nitrites 
are  stable  and  well-characterized  compounds.  When  an  acid  is 
added  to  a  nitrite,  the  nitrous  acid  which  is  formed  at  once  breaks 
down  into  its  anhydride  and  water,  and  the  anhydride  is  further 
decomposed,  so  that  N02  and  NO  are  produced.  By  passing  im- 
pure nitrogen  trioxide  (formed  by  the  reduction  of  nitric  acid)  into 
ice-cold  water,  a  blue  liquid,  which  possibly  is  nitrous  acid,  is  pro- 
duced, but  the  slightest  increase  in  temperature  causes  the  latter  to 
change  into  nitric  acid  and  nitric  oxide. 

We  have  already  studied  the  manner  in  which  nitrites  are 
formed  by  heating  nitrates,  so  nothing  more  need  be  added  except- 
ing the  statement  that  many  nitrites  can  best  be  prepared  from 
potassium  nitrite  by  double  decomposition  (  page  57),  while  potas- 
sium nitrite  is  produced  by  heating  nitrate  of  potassium  with  lead. 

The  hyponitrites  and  hyponitrous  acid  alone  remain  for  dis- 
cussion. 

The  hyponitrite  of  potassium  can  be  prepared  by  reducing  ni- 
trate of  potassium  with  sodium  amalgam*  the  latter  substance, 

*  For  sodium  amalgam  see  foot-note,  reference.  32  of  appendix.  For  the 
preparation  of  hyponitrous  acid  from  hydroxylamin  and  nitrous  acid  see  Wisli- 
cenus,  Ber.  d.  Deutsch.  Chem.  Gesell.  26  ;  771,  and  Paal,  ibid.  1027. 


NITROGEN   OXIDES    AND   ACIDS;    TABLE    OF. 


209 


when  in  contact  with  water,  forms  sodium  hydroxide  and  hydro- 
gen, while  hydrogen  in  the  nascent  state  robs  the  potassium  nitrate 

of  its  oxygen  :  — 

KN03  +  4  H  =  KNT0  +  2  H2  0  . 

The  hyponitrite  of  silver,  which  is  insoluble  in  water,  can  be  pre- 
pared by  adding  silver  nitrate  to  potassium  hyponitrite :  — 

Ag  N03  +  KNO  ==  KN08  +  Ag  NO, 

and  the  free  acid  can  be  formed  from  the  latter  by  the  addition  of 
hydrochloric  acid :  — 

Ag  NO  +  H  Cl  =  Ag  Cl  +  HNO.* 

Hyponitrous  acid  exists  only  in  very  dilute  solutions;  when  warm  3d 
or  when  allowed  to  stand,  it  decomposes,  yielding  nitrous  oxide  (its 
anhydride)  and  water  :  — 

2  HNO  =  H2  0  +  N2  0  . 
The  acid  is  of  no  practical  importance. 

In  the  following  table  the  nitrogen  compounds  are  compared  with 
those  of  chlorine  :  — 


OXIDES. 

ACIDS. 

NAMES. 

OXIDES. 

ACIDS. 

NAMES. 

C12O 

HO   Cl 

Hypochlorous  acid 

N2O 

HO   N* 

Hyponitrous  acid 

C1203 

HO2C1 

Chlorous             " 

N203 

HO2N§ 

Nitrous             « 

(Cl206)t 

HO3C1 

Chloric               " 

N206 

HO3N 

Nitric               " 

(Cl207)t 

H04C1 

Perchloric          " 

*  Silver  chloride  is  insoluble  in  water.  The  molecules  of  hyponitrous  acid 
have,  probably,  a  molecular  weight  which  is  double  that  of  the  formula  weight 
(  Ho  N2  O2  ).  This  conclusion  is  rendered  probable  from  a  study  of  some  of  the 
organic  derivatives  of  hyponitrous  acid.  See  Meyer  and  Jacobson;  Lehrbuch 
der  Organischen  Chemie ;  I.  207. 

t  .Oxides  C1.2  O5 ,  C1.2  O7  do  not  exist,  the  corresponding  acids,  HO3  Cl, 
HO4C1,  do. 

|  Hyponitrous  acid  breaks  down  into  its  anhydride,  N2  O,  and  water,  but 
cannot  be  formed  by  dissolving  N2  O  in  water.  It  probably  has  a  molecular 
weight  represented  by  the  formula  (  HON  )2  . 

§  Nitrous  acid  is  stable  only  in  very  cold  water.  The  existence  of  the 
anhydride  N2  O3  is  doubtful. 

The  salts  of  these  acids  are  much  more  stable  than  the  acids  themselves. 
All  of  the  acids  are  powerful  oxidizers,  all  of  the  oxides  are  unstable.  Those 
of  chlorine  are  explosive,  those  of  nitrogen  support  combustion. 

Cl  O2  and  NO2  are  not  the  anhydrides  of  acids.  The  former  on  addition 
of  water  forms  chloric  and  chlorous  acids,  Cl  O2  +  H2  O  =  H  Cl  O2  +H  Cl  O3; 


210 


NITROGEN   OXIDES;    HEATS    OF   FORMATION. 


The  heats  of  formation  of  the  oxides  of  nitrogen,  as  far  as  they 
have  been  ascertained,  are  given  in  the  following  table  :  — 


N20 

N  0 

N  02 

-180K 
—  210  K 
-    77  K 
131  K 

Nitric  oxide  should  be  less  stable  than  nitrous  oxide  ;  as  a 
consequence,  NO  is  changed  to  N2O  by  moist  iron  filings, 
zinc  dust,  etc.  NO2  is  more  stable  than  NO,  and  is  produced 
therefrom  readily  by  the  addition  of  oxygen.  The  higher  ox- 
ides are  more  stable  than  those  with  less  oxygen.  N2  Os  is 
an  exothermic  compound  ;  it  is  a  crystalline  solid  which  can 
easily  be  formed  from  nitric  acid  by  extracting  water. 

HNO2 
HN03 

308  K* 
491  K* 

Nitric  acid  has  a  greater  heat  of  formation  than  has  ni- 
trous acid;  it  is  therefore  the  acid  of  oxygen  and  nitrogen 
which  is  most  easily  formed.  The  same  rule  is  observed  in 
the  halogen  and  oxygen  families,  where  those  acids  which 
contain  the  most  oxygen  are  the  most  stable. 

the  latter  nitric  acid  and  nitric  oxide,  3  NO2  +  H2  O  =  2  H  NO3  +  NO.     At 
low  temperatures  NO2  becomes  N2  O4 ,  and  probably  Cl  O2  becomes  C12  O4  . 

The  acids  are  all  unibasic. 

*  Acids  in  solution. 


PHOSPHORUS;    OCCURRENCE.  211 


CHAPTER   XXVII. 

PHOSPHORUS  AND  PHOSPHINE. 

Phosphorus  ;  symbol,  P  ;  atomic  weight,  31 ;  specific  gravity  of  yellow 
phosphorus,  1.83,  of  red  phosphorus,  2.1.  Specific  gravity  of 
vapor,  air  =  1,  is  4.16,  H2  =  2,  is  119.80  ;  molecular  weight,  124  ; 
molecule,  P4 .  Phosphine ;  formula,  PH3 ;  specific  weight,  air 
=  1,  is  1.185,  H2  =  2,  i's  34.12  ;  molecular  weight,  34  ;  1  c.c.  o/ 

£A,e  gas  at  0°and  .76  m.  weighs  .0015276  gram. 
PHOSPHORUS  never  occurs,  as  such,  in  nature ;  indeed,  such  a 
possibility  is  precluded  by  the  chemical  nature  of  the  element,  as 
an  example  of  which  we  have  but  to  recall  the  energy  with  which  it 
burns  in  oxygen.  The  compounds  of  phosphorus  which  are  most 
frequently  found  are  :  — 

Apatite,  a  combination  of  calcium  phosphate  and  calcium  chloride  (or 
fluoride),  Ca3  (  PO4  )2 ,  Ca  C12  . 

Osteolite,  calcium  phosphate,  Ca3(  PO4  )2. 

Vimanite,  ferrous  phosphate,  Fe3  (  PO4  )2  +  8  H2  O.  Phosphates  of  alu- 
minium and  of  lead  also  occur. 

Phosphates  are  always  present  in  the  soil ;  they  are  essential  to 
the  growth  of  plants  and  are  taken  up  by  the  roots,  so  that  plant 
ashes,  especially  those  of  the  cereals,  often  contain  large  quantities 
of  the  phosphates  of  calcium  and  magnesium ;  the  latter  Wd  their 
way  into  the  animal  organism  (calcium  phosphate  is  the  most  im- 
portant inorganic  constituent  of  the  bones)  ;  the  waste  products 
are  returned  to  the  soil  by  means  of  the  solid  excrements  and  the 
urine;  so  that,  as  manure,  they  are  once  more  brought  into  the 
proper  condition  to  play  their  part  in  plant  growth.  These  changes 
taking  place  with  phosphoric  dteid  remind  us  forcibly  of  the  similar 
ones  encountered  with  carbon  dioxide  and  ammonia ;  none  of  these 
necessary  substances  are  the  permanent  property  of  any  one  organ- 
ism ;  they  are  simply  borrowed  for  a  time,  and  must  be  returned  to 
the  place  from  which  they  came. 

Phosphorus  was  discovered  by  a  Hamburg  alchemist  named 
Brand,  who  accidentally  prepared  the  element  while  searching  for 


212  PHOSPHORUS;    HISTORY,    PROPERTIES. 

the  philosopher's  stone.  Subsequently  Kunkel  published  an  account 
in  which  he  described  a  method  of  obtaining  the  substance,  but 
until  the  middle  of  the  last  century  the  supply  was  so  small  that 
phosphorus  was  a  very  expensive  article ;  it  was  exclusively  pre- 
pared from  decaying  urine,  and  the  price  in  England  was  ten  ducats 
an  ounce.  At  a  later  date  a  method  was  discovered  by  which  phos- 
phorus could  be  obtained  from  the  calcium  phosphate  procured 
either  from  mineral  deposits  or  from  bones ;  but  even  then  it  was 
mainly  kept  as  a  curiosity  until  the  introduction  of  matches  ren- 
dered its  cheap  production  necessary.  At  the  present  time  phos- 
phorus is  prepared  from  bones  by  first  burning  the  latter  in  order 
to  destroy  the  organic  matter  contained  in  them ;  the  calcium  phos- 
phate is  then  changed  to  the  primary  phosphate  of  calcium 
(Ca  (H2P04)2)  by  means  of  sulphuric  acid.  Primary  phosphates 
are  soluble  in  water,  so  that  a  solution  can  be  formed  which  is 
further  evaporated  and  heated,  by  which  means  the  primary  phos- 
phate of  calcium  loses  water  and  is  converted  into  calcium  meta- 
phosphate  (Ca  (P03  )2 ) ;  and  the  latter  substance,  when  heated 
with  charcoal  and  sand,  yields  phosphorus. 

Phosphorus  exists  in  two  allotropic  forms,*  the  most  common  of 
which  is  a  slightly  yellow,  wax-like  solid,  which  becomes  brittle 
when  cold,  and  which  is  readily  soluble  in  carbon  bisulphide ;  it 
melts  at  44°  and  boils  at  250°,  forming  a  colorless  vapor  which  has 
a  specific  gravity  of  4.16  at  red  heat.  This,  with  hydrogen  as  two, 
gives  119.8,  while  the  molecular  weight  of  P4  would  be  124.  The 
observed  specific  gravity  is  therefore  somewhat  less  than  the  molec- 
ular weight,  124,  a  fact  which  probably  finds  its  explanation  in 
the  decomposition  of  some  of  these  complex  molecules  into  simpler 
ones.  At  white  heat  the  specific  gravity  of  phosphorus  vapor  has 
diminished  to  3.14,  so  that  at  this  temperature  nearly  all  of  the  P4 
molecules  have  dissociated  into  those  having  the  composition  P2 . 

When  a  solution  of  ordinary  phosphorus  in  carbon  bisulphide  is 
exposed  to  the  sunlight,  the  other,  insoluble,  red,  amorphous  modifi- 
cation of  the  element  gradually  separates.  This  change  can  be 
accomplished  more  quickly  and  effectually  by  heating  phosphorus 
to  about  300°  in  a  closed  vessel ; 60  the  same  transformation  also 

*  A  form  of  phosphorus  resembling  flowers  of  sulphur  has  been  prepared 
by  rapidly  cooling  phosphorus  vapors.  This  may  be  a  third  allotropic  form  of 
phosphorus. 


PHOSPHORUS  ;    ALLOTROPIC    FORMS    OF.  213 

occurs  through  the  influence  of  electricity.  Amorphous  phosphorus 
is  a  dark  red  substance  which  is  generally  produced  in  the  form  of 
a  powder,  the  specific  gravity  of  which  is  2.1.  When  heated  to 
a  temperature  of  358°  in  a  vacuum,  and  even  more  rapidly  at  445°, 
amorphous  phosphorus  is  changed  back  to  the  yellow  variety ;  the 
kindling  temperature  of  the  former  about  coincides  with  this  point. 
The  transformation  of  the  element  from  its  ordinary  crystalline 
form  into  the  amorphous  one  is  accomplished  only  when  the  phos- 
phorus is  under  a  pressure  of  several  atmospheres  and  at  a  higher 
temperature. 

Red  phosphorus  is  perfectly  insoluble  in  carbon  bisulphide, 
ether,  and  similar  substances  by  which  the  other  allotropic  form  is 
readily  dissolved.  It  can  be  exposed  to  the  atmosphere  for  any 
length  of  time  without  change* while  the  other  variety  will  absorb 
oxygen,  melt,  and,  under  proper  conditions,  may  take  fire  spontane- 
ously. Yellow  phosphorus,  when  placed  in  warm,  moist  air  and  in 
the  dark,  emits  a  pale  white  light  f  which  is  in  part  due  to  the 
slow  oxidation  of  the  element.  \ 

Yellow  phosphorus  is  an  intense  poison ;  even  small  doses  cause 
local  inflammations  in  various  organs  of  the  body,  and  have  a  sec- 
ondary effect  on  the  nervous  system.  The  serious  symptoms  caused 
by  poisoning  with  phosphorus  first  become  apparent  some  hours 
after  taking ;  they  manifest  themselves  by  intense  pain  in  the  gas- 
tric region,  finally  extending  throughout  the  entire  abdomen;  the 
vomit  will  contain  phosphorus,  have  a  peculiarly  garlic-like  odor, 
and  will  be  luminous  in  the  dark;  the  patient  is  restless,  fearful, 
and  trembling.  The  post-mortem  examination  reveals  inflammation 
of  the  mucous  membrane  of  the  stomach,  accompanied  by  fatty  de- 
generation of  the  liver,  kidneys,  and  heart.  Fatal  doses  are  from 
.2  to  .5  gram.  Cases  of  phosphorus  poisoning  are  not  uncommon, 
as  phosphorus  mixed  in  a  dough  made  of  cold  water  and  flour  is 
frequently  used  as  a  rat-poison ;  this  has  especially  been  the  case 
since  the  element  has  become  quite  cheap  by  reason  of  its  use  in 
the  manufacture  of  matches,  the  heads  of  a  number  of  varieties  of 
which  are  made  of  a  mixture  of  gum  arabic  and  phosphorus. 

Yellow  phosphorus  is  soluble  in  carbon  bisulphide,  ether,  and 

*  This  statement  is  denied  by  Pedler,  Journ.  Chem.  Soc.  1890,  608. 
t  So-called  phosphorescence. 

J  That  this  is  not  entirely  so  is  proven  by  the  fact  that  phosphorus  is  not 
luminous  in  dry  oxygen  below  20°  C.,  or  in  that  gas  under  pressure. 


214  PHOSPHINE;  PREPARATION. 

ethereal  oils ;  it  is  insoluble  in  alcohol  and  water,  but  volatile  in 
the  vapors  of  the  latter.  When  slowly  oxidized  in  moist  air  it 
changes  to  phosphorous  acid  :  — 

2P+3  0  =  P203, 

P203  +  3  H20=2  P03H3, 

this  oxidation  is  supposed  to  be  the  cause  of  the  phosphorescence  of 
the  element.  When  burned,  both  yellow  and  red  phosphorus  yield 
phosphorus  pentoxide,  which  can  further  unite  with  water  to  form 
phosphoric  acid :  — 

2P  +  50  =  P205, 

P2  05  +  3  H2  0  =  2  P04  H3 . 

The  element  will  combine  with  chlorine,  bromine,  or  any  of  the 
halogens,  just  as  it  will  with  oxygen.  (See  page  63.)  Both  of 
the  oxides  of  phosphorus  are  anhydrides,  and  all  of  the  halogen 
compounds  are  decomposed  by  water.  (See  pages  80,  84.) 

Phosphorus  forms  three  compounds  with  hydrogen,  PH3 ,  phos- 
phine,  P2  H4 ,  liquid  hydrogen  phosphide,  and  P4  H2 ,  solid  hydrogen 
phosphide. 

Phosphine  is  a  gas  which  bears  the  same  resemblance  to  ammo- 
nia that  hydrogen  sulphide  does  to  water.  As  ammonia  is  formed 
with  difficulty  by  the  direct  union  of  nitrogen  and  hydrogen,  we 
would  scarcely  expect  phosphine  to  be  produced  in  a  similar  way, 
and  yet  the  compound  seems  to  be  readily  procured  as  a  result  of 
the  action  of  nascent  hydrogen  upon  phosphorus.*  This  unex- 
pected result  is  possibly  due  to  the  fact  that  the  breaking  stress 
of  the  molecules  of  phosphorus  is  less  than  the  same  for  those  of 
nitrogen.  The  best  method  of  preparing  phosphine  for  laboratory 
use  is  by  decomposing  calcium  phosphide  with  water  or  dilute  acids. 
The  formula  of  calcium  phosphide  has  not  been  definitely  ascer- 
tained, but  we  can  compare  this  reaction  with  similar  ones  in  which 
hydrochloric  acid  or  hydrogen  sulphide  has  been  produced  by  the 
action  of  an  acid  upon  a  chloride  or  a  sulphide.  Another  way, 
which  has  less  to  recommend  it,  but  which  is  more  frequently  used, 

*  By  throwing  small  pieces  of  phosphorus  into  a  flask  in  which  zinc  and 
dilute  sulphuric  acid,  or  tin  and  sulphuric  acid,  are  generating  hydrogen,  the 
temperature  being  about  70°.  Compare  J.  Brossler;  Berichte  der  Deutschen 
Chemischen  Gesellschaf t ;  1881,  1757. 


PHOSPHINE;    PROPERTIES.  215 

is  by  heating  small  pieces  of  phosphorus  in  a  solution  of  potassium 
hydroxide.61  * 

The  gas  formed  by  either  of  these  methods  is  a  mixture  of 
hydrogen  compounds  having  the  formulae  of  PH3  and  P2  H4  (unless 
concentrated  hydrochloric  acid  is  used  to  decompose  the  calcium 
phosphide,  in  which  event  the  substance  corresponding  in  structure 
to  ammonia  is  alone  produced).  This  mixture  of  gases  takes  fire 
spontaneously  when  it  comes  in  contact  with  the  air,  while  pure 
phosphine,  PH3,  does  not  possess  this  property.  The  spontane- 
ously inflammable  gas  can  be  altered  in  this  respect  by  passing 
it  through  a  tube  cooled  with  snow  and  salt,  for  by  this  means 
the  liquid  hydrogen  phosphide  (P2H4)  is  condensed,  while  the 
phosphine  passes  on,  to  be  used  as  occasion  requires. 

Phosphine  is  a  colorless  gas  with  an  intensely  disagreeable, 
garlic-like  odor.  Its  specific  gravity,  air  =  1,  is  1.185,  which, 
H  =  2,  is  34.12.  The  molecular  weight  of  PH3  is  therefore 
34.024,  for  analysis  has  proven  that  in  phosphine  there  are  31 
parts  of  phosphorus  and  3.024  of  hydrogen  by  weight.  It  fol- 
lows that  31  represents  the  maximum  value  for  the  atomic  weight 
of  phosphorus,  for,  as  the  molecular  weight  of  phosphine  is  known, 
we  cannot  imagine  any  atomic  weight  for  phosphorus  greater  than 
this  number  without  believing  that  we  have  a  fraction  of  an  atom 
of  phosphorus  in  PH3.  When  phosphine  is  heated,  or  when  elec- 
tric sparks  are  passed  through  it,  the  gas  breaks  down  into  phos- 
phorus and  hydrogen;  in  this  case  two  volumes  of  hydrogen 
phosphide  yield  three  of  hydrogen,  the  phosphorus,  being  solid, 
when  separated  exerts  no  influence  on  the  volume  of  the  gas  as  a 
whole.  (Compare  pages  99  and  138.) 

H 


r  =       TT  +  phosphorus. 

HP1S    H* 

H 

*  The  reaction  is  said  to  take  place  as  follows  :  — 

3  KOH  +  4  P  +  3  H20  =3  KH2PO.2  +  PH3. 

The  primary  hypophosphite  of  potassium  (KPO  +  H.2O)  would  thus  be 
formed.  The  phosphine  generated  always  contains  hydrogen,  so  that  its  for- 
mation is  probably  due  to  that  element  acting  in  the  nascent  state. 


216  PHOSPHINE;    PHOSPHONIUM   COMPOUNDS. 

From  this  equation  it  is  evident  that  two  molecules  of  phosphine 
produce  three  of  hydrogen,  and  the  terms  "  volume "  and  "  mole- 
cule "  can  be  used  interchangeably,  as  we  saw  on  page  70. 

Phosphine  can  be  mixed  with  pure  oxygen  without  taking  fire, 
but  if  the  pressure  is  suddenly  diminished  the  gases  will  explode. 
In  the  air  the  kindling  temperature  is  149°,*  the  products  of  the 
combustion  are  phosphoric  anhydride  and  water :  — 

2PH3  + 80  =  P205+3H20, 

and  these  two  substances  naturally  combine  to  form  phosphoric 
acid :  — 

P20&  +  3HaO  =  2H3P04. 

Of  course,  chlorine,  bromine,  or  iodine  would  act  on  phosphine  in 
a  manner  parallel  to  the  action  of  oxygen,f  the  products  of  the 
reaction  being  halhydric  acids  and  the  corresponding  halogen  com- 
pounds of  phosphorus ;  for  example  :  — 

PH3  +  6C1  =  PC13  +  3HC1. 

So  readily  is  phosphine  decomposed,  that  even  sulphur  when 
placed  in  contact  with  that  substance,  changes  it  to  the  sulphide 
of  phosphorus  and  hydrogen  sulphide,  so  that  here  we  have  an 
instance  in  which  sulphur  causes  changes  similar  to  those  produced 
by  oxygen  and  the  halogens.  It  need  scarcely  be  added  that  con- 
centrated sulphuric  or  nitric  acid  will  decompose  phosphine  just  as 
they  would  hydrobromic  acid,  hydroiodic  acid,  or  sulphuretted  hy- 
drogen. When  phosphine  is  passed  into  solutions  of  metallic  salts 
it,  in  many  cases,  produces  the  corresponding  phosphides  of  the 
metals  (compare  sulphuretted  hydrogen,  page  99).  Phosphine  is 
sparingly  soluble  in  water,  and  is  very  poisonous,  so  that  all  opera- 
tions in  which  it  is  generated  must  be  conducted  either  in  the  open 
air,  or  under  a  hood  with  a  strong  draught.  The  gas  changes  to  a 
liquid  at  —  85°  and  becomes  solid  at  —  133°. 

Phosphine  can  unite  with  the  halhydric  acids  to  form  phospho- 
nium  compounds  exactly  as  ammonia  does  in  the  production  of 
ammonium  salts,  the  group  PH4  being  termed  phosphonium  for  the 
same  reason  that  NH4  is  ammonium  :  — 

*  The  friction  of  the  glass-stopper  in  the  neck  of  a  bottle  filled  with  phos- 
phine may  be  sufficient  to  ignite  the  gas. 

t  Chlorine  and  phosphine,  when  mixed,  explode  very  violently. 


LIQUID   HYDROGEN   PHOSPHIDE.  217 

PH3  +  HI  =  PH4  1,  phosphonium  iodide, 
NH3  +  HI  =  NH4  1,  ammonium  iodide. 

As  phosphine  is  readily  oxidized  and  is  much  less  basic  in  its 
character  than  is  ammonia,*  it  will  not  unite  with  acids  containing 
oxygen  ;  in  this  respect  it  differs  from  ammonia.  The  phosphonium 
compounds  are  readily  decomposed  by  water  or  by  alkalies  :  t  — 

PH4I  +  H20  =  PH3  +  HI+H20, 
PH4I  +  KOH  =  PH3  +  KI  +  H20, 

the  latter  reaction  being  exactly  like  those  observed  with  ammonium 

salts  :  — 

+  KOH  =NH3  +  KI  +  H20. 


It  follows  from  the  above  that  phosphonium  salts  cannot  be  formed 
where  water  is  present. 

The  compound  P2  H4  is  a  liquid  at  ordinary  temperatures.  It  is 
isolated,  as  was  stated  above,  by  cooling  the  mixture  of  gases 
obtained  by  one  of  the  ordinary  methods  in  use  for  the  production 
of  phosphine.  It  is  a  colorless,  highly  refractive  liquid,  which  boils 
at  about  35°,  and  which  takes  fire  spontaneously  when  exposed  to 
the  air.  The  determination  of  the  specific  gravity  of  the  vapor 
shows  it  to  have  a  molecular  weight  of  66  ;  its  formula  is,  there- 
fore, PH2  —  PH2  ;  the  analogous  compound  of  nitrogen  is  hydra- 
zin,  NH2  —  NH2;  but,  unlike  the  latter,  liquid  hydrogen  phosphide 
is  not  basic,  and  can,  therefore,  form  no  salts.  This  fact  is  not  sur- 
prising if  the  same  diminution  of  basic  properties  takes  place  with 
the  hydrogen  compounds  of  phosphorus,  as  was  observed  in  the  case 
of  the  similar  ones  containing  nitrogen.  (  Page  193,  foot-note.) 

A  solid  compound,  P2  H,  is  formed  by  treating  phosphine  with 
chlorine  which  has  been  highly  diluted  with  carbon  dioxide.  By 
this  means  a  part  of  the  hydrogen  of  phosphine  is  removed,  and  a 
yellow  powder  is  produced,  which,  when  dry,  can  be  heated  as  high 
as  150°  without  taking  fire. 

*  See  page  177. 

t  Phosphonium  bromide  or  phosphonium  iodide  is  more  easily  formed 
than  is  phosphonium  chloride.  The  latter  salt  is  prepared  by  subjecting  a 
mixture  of  equal  volumes  of  phosphine  and  of  hydrochloric  acid  to  a  pressure 
of  20  atmospheres;  phosphonium  iodide  is  produced  by  direct  union  at  ordinary 
pressure. 


218 


HYDROGEN   PHOSPHIDES;    TABLE   OF. 


The  following  table  shows  the  relationship  between  the  com- 
pounds just  discussed  and  the  corresponding  ones  containing 
nitrogen :  — 


NH3  ,  ammonia, 
N2  H4  ,  hydrazin, 

N,  H,  azoimid. 

PH3,  phosphine. 
P2  H4  ,      liquid      hydrogen  : 
phosphide. 

NH3  +  HX  =  NH4  X,  Ammonium  salts. 
PH3  +  HX  =  PH4X,  Phosphonium  salts. 

Phosohonium  fluoride  has  not.  as  vet. 

PaH,  solid  hydrogen  phos- 
phide. 


been  prepared;  with  this  exception 
X  represents  any  halogen. 


The  compounds  of  nitrogen  are  all  gases;  they  are  not  spontaneously 
inflammable,  while  those  of  phosphorus  are  either  gaseous,  liquid,  or  solid, 
and  burn  with  the  greatest  ease. 


PHOSPHORUS;    HALIDES   OF. 


219 


CHAPTEE   XXVIII. 


THE  COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS, 
AND  WITH  OXYGEN  AND  THE  HALOGENS. 

As  the  atomic  weights  in  this  family  increase,  an  increasing 
stability  of  the  compounds  formed  with  the  halogens  is  observed. 
Those  of  nitrogen  are  very  explosive  substances  ;  but  in  the  case 
of  the  element  under  consideration  a  number  of  quite  stable  chlo- 
rides, bromides,  and  iodides  have  been  accurately  studied  ;  indeed, 
some  of  these  can  be  classed  among  our  most  important  laboratory 
reagents.  They  are  given  in  the  following  table  :  — 


FLUORIDES. 

CHLORIDES. 

BROMIDES. 

IODIDES. 

PF3,  phosphorus  trifluoride, 
PF8,  phosphorus    pentafluo- 
ride, 
An  iodide  of  phosphorus,  PI 

P  C13,  trichloride, 
P  Cls,  pentachloride, 

,  phosphorus  di-iodide, 

P  Br3,  tribromide, 
P  Br6,  pentabro- 
mide, 

exists. 

PI3,  triiodide. 

The  above  compounds  are  all  substances  which,  because  they 
are  the  halides  of  a  not-rnetal,  are  readily  decomposed  by  water  to 
form  the  corresponding  acid  of  phosphorus,  together  with  the 
hydrogen  compound  of  the  halogen  which  was  united  with  that 
element.  We  have  seen  that  this  instability  in  the  presence  of 
water  has  been  made  use  of  in  the  preparation  of  hydrobromic  and 
hydroiodic  acids.  (Pages  80  and  85.) 

The  character  of  the  trihalogen  compounds  changes  somewhat 
writh  the  nature  of  the  halogen,  the  boiling  point  increases  with  the 
increase  of  the  molecular  weight,  just  as  it  does  in  the  case  of  the 
free  elements,  while  tne  readiness  with  which  these  substances  are 
decomposed  is  also,  apparently,  greater  in  the  bromide  and  iodide 
than  it  is  in  the  fluoride  and  chloride. 

P  F3  is  a  gas,  liquid  at  —  10°  under  a  pressure  of  40  atmospheres. 
P  C13  is  a  liquid  which  boils  at  76°,  heat  of  formation  755  K. 
P  Br3  is  a  liquid  which  boils  at  175°,  heat  of  formation  448  K. 
P I3  is  a  solid  which  melts  at  55°  and  is  decomposed  by  boiling. 
PF3,  specific   gravity  of  vapor,  air  =  1,  is  3.02,  which,  H2  =  2,  is  88;    the 
molecular  weight  is  88. 


220  PHOSPHORUS  ;   TRIHALIDES  J   PENTAHALIDES. 

P  C13,  specific  gravity  of  vapor,  air  =  1,  is  4.8,  which,  H2  =  2,  is  138;  the 
molecular  weight  is  137.35. 

PBr3,  specific  gravity  of  vapor,  air  =  1,  is  9.7,  which,  H2=  2,  is  279.3;  the 
molecular  weight  is  270.88. 

From  these  determinations  of  the  specific  gravities  of  the  vapors  it 
is  evident  that  the  general  formula  of  all  of  the  trihalides  of  phos- 
phorus is  PX3,  so  that  phosphorus  is  trivalent  in  these  compounds, 
just  as  it  is  in  phosphine,  or  just  as  nitrogen  is  in  ammonia.  The 
trihalides  of  phosphorus  are  all  formed  by  treating  phosphorus  with 
an  amount  of  halogen  insufficient  to  produce  the  compounds  PX5,  * 
and  when  they  are  decomposed  by  water  they  break  down  as 
follows :  — 

X  +  HOH  f  OH 

X  +  HOH  P  •]  OH  +  3  HX, 

X  +  HOH  (  OH 

so  that  phosphorous  acid  results  in  all  cases. 

Phosphorus  is  unsaturated  in  the  trihalogen  compounds,  and  it 
is  therefore  capable  of  a  further  addition  to  form  pentahalides,  the 
valence  of  the  element  increasing  from  three  to  five :  — 
PX3  +  2  X  =  PX5 . 

P  F6  is  a  gas  which  liquefies  at  16°,  46  atmospheres  pressure. 

P  C16  is  a  solid  which  melts  at  148°  (under  diminished  pressure),  and  which 
boils  at  about  160°  —  165°,  with  partial  decomposition. 

P  Br5  is  a  solid  which  decomposes  into  bromine  and  phosphorus  tribromide 
.at  100°. 

These  compounds,  when  added  to  water,  yield  phosphoric  acid, 
as  will  be  seen  from  the  following  :  — 

X  +  HOH  r_0  — H 

X  +  HOH  __0  — H 

X  +  HOH      =       P^l— 0  —  H+5HX. 

X  +  HOH  |_0  — H 

X  +  HOH  [— 0  —  H 

From  the  above  equation  we  would  expect  the  production  of  normal 
phosphoric  acid,  but,  as  we  have  already  seen,  the  normal  acids 
have  the  greatest  tendency  to  separate  water,  by  this  means  yield- 
ing more  stable  compounds  (pages  130  and  131),  so  that  P  (OH)6 
breaks  down  to  form  ordinary  phosphoric  acid :  — 

*  Excepting  the  trifluoride  which  is  formed  by  a  somewhat  complicated 
process. 


PHOSPHORUS;    OXY-HALIDES    OF.  221 

P(OH)5  =  P04H3  +  H20. 

The  difference  between  the  tri-  and  pentachloride  of  phosphorus 
lies  in  the  different  amounts  of  chlorine  as  compared  to  the 
quantity  of  phosphorus  contained  in  each  ;  in  the  trichloride 
phosphorus  is  trivalent,  in  the  pentachloride  it  is  quinquivalent; 
the  one  yields  phosphorous  acid,  the  other  phosphoric  acid  by 
the  addition  of  water,  so  that,  plainly,  the  same  difference  in  the 
valence  of  phosphorus  exists  in  these  acids  as  is  found  to  exist 
in  the  valence  of  the  element  in  the  two  chlorides. 

Compounds  containing  both  bromine  and  chlorine  can  be  formed 
by  adding  bromine  to  the  trichloride,  or  chlorine  to  the  tribromide 
of  phosphorus.  Under  proper  conditions  a  portion  of  the  chlorine 
or  bromine  in  the  pentachloride  or  bromide  of  phosphorus  can  be 
replaced  by  oxygen  ;  the  result  is  the  production  of  an  oxy-chloride 
or  bromide  of  phosphorus.  Phosphorus  oxy-chloride  has  a  chemical 
character  which  is  analogous  to  that  of  the  similar  compounds  of 
sulphur  discussed  011  page  145. 

POC13  —  phosphorus  oxychloride,  liquid,  boils  at  110°. 

PO  Br3  —  phosphorus  oxybromide,  solid,  melts  at  55°,  and  boils  at  193°. 

These  substances  are  produced  by  adding  the  calculated  amount 
of  water  to  the  pentachloride  or  bromide  of  phosphorus  :  — 

PX5  +  H2  0  =  POX3  +  2  HX. 

In  effecting  this  change,  one  atom  of  oxygen  has  taken  the 
place  of  two  atoms  of  chlorine  or  bromine,  so  that  the  constitution 
of  these  compounds  is  as  follows  :  — 


-x 


p- 

n- 


t-x 

and  this  structural  formula  is  further  supported  by  the  vapor  densi- 
ties of  the  oxychloride  and  oxybromide,  which  exactly  agree  with 
the  theory. 

Both  phosphorus  oxychloride  and  bromide  are  converted  into 
ordinary  phosphoric  acid  by  the  addition  of  water,  therefore  the 
latter  compound  contains  three  hydroxyl  groups  and  is  formed  as 
follows  :  — 


222  PHOSPHORUS   OXYCHLOKIDE. 

=o  r=o 

-Cl  +  HOH  I  -OH 


p      -  _   p     - 

-Cl  +  HOH  —OH" 

Cl  +  HOH  L—  OH 

All  of  the  halogen  compounds  of  phosphorus  fume  in  the  air, 
because  moisture  decomposes  them  while  liberating  hydrogen  chlo- 
ride, bromide,  or  iodide.  Phosphorus  oxychloride  and  oxybromide 
have  an  indescribably  unpleasant  odor,  so  that  all  work  with  these 
substances,  as  well  as  with  the  pentahalogen  compounds,  must  be 
so  conducted  that  the  vapors  cannot  be  inhaled. 


PHOSPHOKUS  ;    TKIOXIDE.  223 


CHAPTER   XXIX. 

THE  COMPOUNDS  OF  PHOSPHORUS  WITH  OXYGEN,  AND  WITH 
OXYGEN  AND  HYDROGEN. 

PHOSPHORUS  forms  four  oxides ;  two,  P2  08  and  P2  05 ,  correspond 
to  nitrogen  trioxide,  N203,  and  nitrogen  pentoxide,  N206,  one,  P02, 
is  analogous  to  N02 ,  and  the  fourth  oxide  probably  has  the  for- 
mula P40.  Some  other  oxides  of  phosphorus  have  been  described, 
but  further  investigation  must  establish  their  identity.  Phosphorus 
trioxide  and  pentoxide  are  both  acidic  anhydrides,  the  one  of  phos- 
phorous, the  other  of  phosphoric  acid.  The  most  common  forms  of 
these  acids  differ  from  the  corresponding  ones  of  nitrogen  by  being 
hydrated,  so  that  phosphorous  acid  is  not  HP02  but  H3  P03 ,  and 
phosphoric  acid  not  HP03  but  H3  P04 . 

Phosphorus  trioxide  is  produced  by  slowly  oxidizing  phosphorus 
in  a  stream  of  oxygen  diluted  with  carbon  dioxide.  It  is  a  crystal- 
line solid  which  melts  at  22°.5  and  which  boils  at  173°. 3,  being 
changed  to  a  colorless  vapor  which  has  a  specific  gravity,  air  =  1, 
of  7.6;  this  indicates  a  molecule  of  the  formula  P406,  as  the  spe- 
cific gravity  calculated  for  P2  03  is  3.8.*  The  oxide  is  completely 
decomposed  when  heated  to  300°,  at  which  temperature  phosphorus 
and  the  oxide  P2  04  are  produced.  Phosphorus  trioxide  is  oxidized 
when  brought  in  contact  with  oxygen ;  the  action  may  even  become 
so  violent  as  to  cause  spontaneous  combustion  to  ensue ;  f  the  oxi- 
dation product  is  phosphorus  pentoxide  ;  the  latter  substance,  as  we 
have  seen,  is  also  produced  when  phosphorus  is  burned  in  air  or 
oxygen  (page  22). 62  Phosphorus  pentoxide  is  a  flaky,  not  crystal- 
line powder ;  under  certain  circumstances  it  can  be  obtained  in  a 

*  Thorpe  and  Tutton,  Journ.  Qhem.  Soc. ;  1890 ;  545. 

t  This  phenomenon  is  possibly  due  to  the  fact  that  a  mixture  of  phosphorus 
trioxide,  pentoxide,  and  phosphorus  is  present  after  slow  oxidation  of  phos- 
phorus ;  the  phosphorus  would  then  cause  the  spontaneous  combustion  of  the 
mass.  Pure  phosphorus  trioxide  unites  with  oxygen  and  becomes  luminous 
when  placed  in  the  gas  and  under  diminished  pressure ;  the  glowing  ceases 
when  the  pressure  is  increased.  This  fact  reminds  us  of  the  similar  phenom- 
enon observed  with  phosphorus  and  oxygen. 


224  PHOSPHOROUS   ACID. 

crystalline  form.  Phosphorus  pentoxide  greedily  absorbs  moisture 
from  the  air;  it  is  therefore  deliquescent;  its  tendency  to  unite 
with  water  is  so  great  that,  if  a  little  of  it  is  placed  in  that  liquid, 
it  dissolves  with  a  hissing  noise  similar  to  that  which  is  heard  when 
a  red-hot  iron  is  immersed.  Phosphoric  anhydride,  because  it  is 
able  to  absorb  perfectly  all  moisture  from  gases,  is  a  favorite  labora- 
tory reagent  for  drying.  It  is  the  more  valuable,  because  the  usual 
drying  agent,  calcium  chloride,*  does  not  completely  remove  water 
from  substances  with  which  it  is  brought  in  contact.  Phosphorus 
pentoxide  is  quite  volatile  if  it  is  carefully  heated  to  250° ;  but 
above  that  temperature  it  changes  into  another,  so-called  poly- 
meric f  form,  which  evaporates  very  slowly  below  bright  red  heat. 

Phosphorous  acid  is  produced  when  the  trioxide  or  trichloride 
is  dissolved  in  water ;  it  is  analogous  to  nitrous  acid,  although  it  is 
much  more  stable.  When  phosphorus  trioxide  unites  with  water 
the  first  product  which  we  should  expect  would  be  H  P02 ,  f  or  :  — 

P2  03  +  H2  0  =  2  H  P02  (see  page  117). 

The  compound  so  produced,  however,  takes  up  one  more  molecule 
of  water  to  form  the  hydrated  acid  H3  P03 :  — 

HP02  +  H20  =  H3P03, 

in  which  condition  only,  the  acid  is  capable  of  existence. 

Phosphorous  acid  contains  three  hydrogen  atoms  in  each  for- 
mula weight,  but  no  more  than  two  of  these  can  be  replaced  by 
metals  at  the  same  time.  The  following  explanation  of  this  phe- 
nomenon seems  the  most  reasonable.  The  character  of  any  chemi- 
cal compound  is  influenced  by  all  of  the  elements  in  that  compound ; 
no  one  element  or  group  of  elements  is  able  entirely  to  suppress 
any  one  of  the  others  with  which  it  is  united.  When  a  hydroxide 

*  It  has  been  shown  that  a  glass  tube  four  inches  in  length,  filled  with 
phosphorus  pentoxide,  will  entirely  dry  a  gas  which  is  slowly  passing  through. 

t  The  polymeric  form  of  a  substance  is  supposed  to  be  produced  by  the 
union  of  simpler  molecules  of  that  substance  to  form  a  more  complicated 
molecule.  Thus,  ordinary  P2  O6 ,  let  us  suppose,  is  formed  of  molecules  each 
of  which  is  composed  of  x  times  the  formula  weight  P2O5,  or  x  (P2O5),  each 
molecule  of  the  polymeric  form  would  then  contain  a  number  of  these  simpler 
molecules,  or  n  (x[P2O5J).  Such  polymeric  forms  are  quite  frequently  met 
with  in  organic  chemistry,  and,  possibly,  the  phenomenon  of  allotropism  may 
in  many  cases  be  caused  by  a  union  of  simpler  molecules  to  form  more  com- 
plex ones. 


PHOSPHITES.  225 

(for  instance,  that  of  potassium)  reacts  with  phosphorous  acid  it 
is  to  be  presumed  that  the  first  product  will  be  the  primary  salt 

(pae-e  140)  :  — 

OH  +  KOH  (  OK 

OH  =  PJOH  +  H20. 

OH  (OH 

The  metallic  element  which  is  present  in  the  salt  after  this  reac- 
tion, renders  the  whole  compound  less  negative,  so  that  the  next 
reaction  :  —  /-  0K  OK 


P 
1 


(  OH  (OH 

would  take  place  less  readily  than  the  first.  The  secondary  salt  has 
now  entirely  lost  all  acid  properties,  owing  to  the  increased  mass  of 
metal  present,  so  that  all  attempts  to  replace  the  third  hydrogen 
atom  will  fail.*  Many  chemists  think  that  experimental  evidence 
has  proven  the  formula  of  phosphorous  acid  to  be  :  — 

v  in 

=  0  _  /  OH 

H  _P  —  OH  and  not  P  —  OH 

—  OH  \  OH 

so  that  it  would  contain  only  two  hydroxyl  groups.  One  of  the 
hydrogen  atoms  would  then  be  joined  to  phosphorus,  thus  rendering 
the  not-metal  quinquivalent  ;  by  means  of  this  hypothesis  they  have 
sought  to  explain  the  fact  that  phosphorous  acid  will  only  form  pri- 
mary and  secondary  salts.  The  remarks  on  page  158  will  apply 
equally  well  in  this  case. 

Phosphorous  acid,  like  sulphurous  and  nitrous  acid,  is  easily 
oxidized,  and  when  so  acted  on  it  forms  phosphoric  acid.  Chlorine, 
bromine,  iodine,  nitric  acid,  and  even  sulphurous  acid  can  bring 
about  this  oxidation.  The  following  equations  will  serve  as  illus- 
trations (see  page  139)  :  — 

a.  2C1      +  H20=2HC1  +  0, 

b.  H3P03  +  0       =H3P04. 
Combining  a  and  b,  we  have  :  — 

c.  H3P03  +  2  Cl  +  H20  =  H3P04  +  2  HCl.f 

*  A  number  of  polybasic  acids  which  are  encountered  in  organic  chemis- 
try show  this  same  character;  they  present  the  phenomenon  of  having  one 
atom  of  the  metal  in  their  salts  more  reactive  than  the  others. 

t  The  pupil  should  practise  the  writing  of  a  large  number  of  equations  in 


226  pHOSPHonic  ACIDS. 

Phosphorous  acid  follows  out  the  general  rule  which  we  ob- 
served to  be  in  force  with  the  chlorine  and  sulphur  acids ;  namely, 
it  changes  into  the  acid  with  greater  amount  of  oxygen  when  it  is 
heated.  The  change  can  be  represented  by  the  following  equa- 

tion  :  -  4  H3P03  =  3  H3P04  +  PH3. 

The  evolution  of  phosphine,  if  the  phosphorous  acid  is  heated  too 
rapidly,  may  sometimes  become  so  violent  that  an  explosion  takes 
place.  Phosphorous  acid  is  a  colorless,  crystalline  solid  which 
melts  at  70°  to  74°,  and  which  is  very  soluble  in  water. 

When  phosphoric  anhydride  is  exposed  to  the  air,  it  deliquesces 
and  is  converted  into  a  phosphoric  acid,  which,  however,  is  not  the 
one  usually  encountered,  but  is  the  less  hydrated  acid  correspond- 
ing to  nitric  or  chloric  acid ;  the  following  will  make  the  parallel- 
ism clear :  — 

P205  +  H20=2HP03, 

Phosphoric  anhydride  -1-  Water  =  Phosphoric  acid. 

N205  +H20  =  2HN03, 

Nitric  anhydride  +  Water  =  Nitric  acid. 

C1205  +H20=2  HC103, 

Chloric  anhydride         +  Water  =  Chloric  acid. 

When  the  phosphoric  acid  so  obtained  is  dissolved  in  an  excess  of 
water  and  allowed  to  stand  for  some  time,  or  when  it  is  boiled  with 
water,  it  takes  up  more  of  that  liquid  to  produce  the  ordinary  form 
of  the  acid ;  in  effecting  this  change,  one  oxygen  atom  together  with 
the  elements  of  water  forms  two  hydroxyl  groups :  — 

(=o          -  ho 

=  0  +  HOH     -p    _g_S 

<_o  — H  _o  — H 


The  acid  with  the  formula  HP03  is  termed  metaphosphoric,  while 
H3  P04  has  the  name  of  orthophosphoric  acid.  Nitric  acid,  HK03 , 
and  chloric  acid,  H  Cl  03 ,  are  therefore  really  meta  nitric  and  meta 

which  cases  of  oxidation  and  reduction  occur.  In  all  cases  he  must  consider, 
first,  the  substance  to  be  oxidized;  second,  the  amount  of  oxygen  which  it 
will  take  up;  third,  the  oxidizing  agent;  and,  fourth,  the  amount  of  oxygen 
which  the  latter  will  yield  and  the  products  which  it  forms  when  it  oxidizes. 
Examples  of  oxidation  have  been  frequently  given  on  the  previous  pages  of 
this  book. 


PHOSPHORIC   ACIDS  :   NOMENCLATURE. 


22-7 


chloric  acids,  but  as  the  corresponding  ortho  acids  (H3!N"04  and 
H3  Cl  04),  are  not  known,  there  is  no  necessity  of  applying  a  special 
designation  to  the  less  hydrated  and  common  form  of  these  sub- 
stances. 

The  nomenclatwe  which  is  applied  to  the  acids  derived  from 
the  other  elements  of  the  nitrogen  family  corresponds  to  that  used 
with  phosphorus,  as  the  following  table  will  demonstrate :  — 


FORMULA. 

NOMENCLATURE. 

H  As02 

Meta-arsenious  acid. 

H3As03 

Ortho-arsenious  acid. 

H  As03 

Meta-arsenic  acid. 

H3As04 

Ortho-arsenic  acid. 

A  similar  system  of  nomenclature  is  also  frequently  applied  for  the 
designation  of  acids  derived  from  elements  belonging  to  other 
groups  than  that  of  nitrogen,  for  instance  :  — 

H2  Si  03  is  the  formula  of  meta  silicic  acid, 
H4Si04  «    "          "         "   ortho      "        «    . 

Unfortunately,  this  system  is  not  rigidly  carried  out  with  all  of  the 
various  hydrated  acids  which  have  been  discovered,  and,  further- 
more, acids  have  been  distinguished  by  the  prefixes  ortho  and  meta 
when  the  difference  between  them  has  nothing  whatever  to  do  with 
their  hydration,  so  that  these  distinctive  names  can  only  be  used 
in  cases  where  they  have  the  sanction  of  custom. 

Orthophosphoric  acid  changes  into  metaphosphoric  acid  at  a  red 
heat,  and  if  this  temperature  is  maintained  for  a  sufficient  length 
of  time,  the  latter  substance  will  finally  evaporate.  When  ortho- 
phosphoric  is  changed  into  metaphosphoric  acid,  a  third  acid  is  first 
produced  as  an  intermediary  product.  This  acid,  having  its  place 
between  the  other  two,  is  (owing  to  the  fact  that  it  is  produced  by 
heating)  termed  pyrophosphoric  acid.  The  changes  which  ortho- 
phosphoric  acid  undergoes  can  be  represented  by  the  following :  — 

2  H3  P  04  =     H4  P2  07     +  H2  0;  and 


Orthophosphoric 
acid. 

H,  P2  0, 

Pyrophosphoric 
acid. 


Pyrophosphoric 
acid. 

=     2  H  P03     +  H2  0. 

Metaphosphoric 
acid. 


228  METAPHOSPHORIC   ACID. 

The  structural  formulae,  representing  the   formation  of  pyrophos- 
phoric  acid,  are  as  follows :  — 


P 


0  0 
OH  HO 
OH  HO 


P=P 


0  0 

OH   HO 
OH   HO 

—  0  — 


Por,2H3P04=H4P207+H20. 


I 1  OHH  10 

n//0 

It  is  supposed  that  two  univalent  groups  —    P —  OH  are  united  by 

\OH 

means  of  oxygen,  just  as  are  two  similar  ones  in  disulphuric  acid 
(described  on  page  154)  ;  such  complicated  acids,  produced  by  the 
separation  of  water  and  the  joining  of  more  complicated  groups 
of  elements  by  means  of  oxygen  atoms,  are  of  quite  frequent 
occurrence.* 

Metaphosphoric  acid  is  a  colorless,  glass-like  substance  (acidum 
phosphoricum  glaciale)  which  greedily  absorbs  moisture  from  the 
air.  It  is  a  monobasic  acid,  and,  according  to  the  rules  which  have 
been  dwelt  upon  in  the  former  portions  of  this  book,  should  only 
form  one  class  of  salts.  This  is  not  the  case,  however,  as  a  number 
of  metaphosphates  of  the  same  metal  are  known ;  the  existence  of 
these  various  metaphosphates  is  explained  by  the  theory  that  meta- 
phosphoric  acid  itself  can  exist  in  several  polymeric  forms.  (See 
page  224.)  f  Metaphosphoric  acid  precipitates  egg  albumen  from 
its  solutions,  a  property  which  distinguishes  it  from  ortho  and  pyro- 

*  From  its  formation  by  heating  pyrophosphoric  acid,  it  seems  likely  that 
ordinary  metaphosphoric  acid  has  the  formula :  — 

r=o  o=n 

r\ 


p    -Q_HOH-O- ip 

I o j 


or,  H2  P2  O6  =  2  HPO3  ;  for  any  other  formula  could  be  produced  only  if  we 
suppose  the  pyrophosphoric  acid  to  be  split  asunder  by  the  very  means  which 
generally  unites  groups  of  atoms,  namely,  by  the  separation  of  water  between 
two  hydroxyl  groups.  Of  course,  this  is  merely  speculation,  but  the  existence 
of  more  than  one  sodium  metaphosphate  seems  to  warrant  the  belief  that  meta- 
phosphoric acid  can  appear  in  polymeric  forms. 

t  The  metaphosphoric  acids  are  supposed  to  have  the  formulae :  — 

HP03 

2(HP03)=H2P206, 
3(HP03)=H3P309, 
or,  in  general,  to  be  n  (HPO3). 


OKTHOPHOSPHATES.  229 

phosphoric  acid,  and  it  changes  to  orthophosphoric  acid  when  it  is 
dissolved  in  water  and  allowed  to  stand ;  it  is  rapidly  converted  by 
boiling  the  solution. 

Orthophosphoric  acid  is  tribasic ;  it  forms  primary,  secondary, 
and  tertiary  salts.  Representatives  of  each  of  these  classes  are 
known,  and  all  are  of  importance.  If  M'  represents  any  univalent 
metal,  then  :  — 

M'  H2  P04  is  the  primary  phosphate, 

M'2  HP04  is  the  secondary  phosphate,  and 

M'3  P04  is  the  tertiary  phosphate ; 
and  if  M"  is  any  bivalent  metal,  then  :  — 

M"  (  H2  P04  )2  is  the  primary  phosphate, 

M'v  (  HP04  )  is  the  secondary  phosphate,  and 

M"3  (  P04  )2  is  the  tertiary  phosphate. 

The  formulae  of  the  last  three  salts  are  self-evident  if  we  consider 
that  a  bivalent  metal  replaces  two  hydrogen  atoms  of  an  acid,  while 
a  primary  salt  has  one,  a  secondary  two,  and  a  tertiary  three  such, 
atoms  replaced  by  that  metal. 

When  the  secondary  phosphates  are  heated  to  a  sufficiently  high 
temperature,  the  pyrophosphate  very  frequently  results,  thus :  — 

2  Na2  HP04  =  Na4  P2  07  +  H2  0, 

and  when  the  primary  ones  are  similarly  treated  the  metaphosphate 
is  produced  :  -  Na  HS  PQ4  =  Na  Po3  +  H2  0. 

All  secondary  and  tertiary  phosphates,  excepting  those  of  the  alkali 
metals  and  of  ammonium,  are  insoluble  in  water,  while,  on  the  other 
hand,  all  primary  phosphates  are  soluble.  The  tertiary  and  second- 
ary phosphates,  therefore,  are  dissolved  by  acids,  because  the  latter 
produce  the  primary  phosphates,  thus  :  — 

Ca3  (  P04  )„  +  4  H  Cl  =  Ca  (  H2  P04  )2  +  2  Ca  C12  ,* 

(  Insoluble  tertiarv  )    ,       •  -,        (  soluble  primary       )    .       -,, 
4      -i  •          -I        i    ",     >  4-  acid  =  •<      ,  .       r  T        r        >•  -4-  salt. 
(  calcium  phosphate  )    '  (  calcium  phosphate  )    ' 

By  the  addition  of  soluble  bases  to  the  soluble  primary  phosphates, 

the  insoluble  tertiary  ones  are  precipated :  — 

3  Ca  (H2P04)2  + 12 NH4OH  =Ca3  (P04)2+  4  (NH4)3  P04  + 12  H2 0.  f 

*  See  pages  140  and  153. 

t  The  formulation  of  a  reaction  such  as  this  is  really  not  as  formidable  as 


230  OKTHOPHOSPHATES. 

The  soluble  secondary  phosphates  of  the  alkali  metals  are  the 
salts  of  those  elements  which  are  usually  encountered ;  the  insoluble 
secondary  phosphates  of  pronounced  metals  like  calcium,  barium, 
or  strontium,  which  form  such  salts,  are  readily  produced  from 
these  by  precipitation :  — 

Na2  HP04  +  Ca  C12  =  Ca  HP04  +  2  Na  Cl,* 

Soluble.  Soluble.      Insoluble.        Soluble. 

The  soluble  tertiary  phosphates  of  the  alkali  metals  have  a 
strongly  alkaline  reaction,f  and  show  the  greatest  tendency  to 
separate  a  portion  of  the  metal  contained  in  them  as  the  metallic 
hydroxide :  — 

Na3  P04  +  H2  0  =  Na2  HPO4  +  Na  OH. 

This  characteristic  bears  out  the  theory  which  we  developed  when 
discussing  phosphorous  acid  (pages  224,  225).  Phosphoric  acid, 
being  a  stronger  acid  than  phosphorous  acid,  will  have  a  greater 
tendency  to  form  sodium  salts,  so  that,  while  the  tertiary  sodium 
phosphate,  although  unstable,  really  does  exist,  the  tertiary  phos- 
phite is  not  known.  Only  primary  and  secondary  salts  of  very  pro- 
nounced metals,  such  as  sodium,  potassium,  and  calcium,  are  known ; 

at  first  glance  it  would  appear  to  be.  The  process  consists  in  simply  neutral- 
izing the  excess  of  acid  hydrogen  atoms  in  the  primary  salt  by  means  of 
ammonia.  As  the  tertiary  phosphate  of  calcium  is  always  formed,  we  must 
necessarily,  in  endeavoring  to  picture  by  atomic  symbols  the  changes  which 
really  take  place,  use  three  times  the  formula  weight  of  calcium  primary  phos- 
phate, because  there  are  three  atoms  of  calcium  in  Ca3  (PO4)2 .  The  remainder 
of  the  work  simply  consists  in  counting  up  how  much  ammonia  it  will  take 
to  replace  all  of  the  remaining  hydrogen  atoms  by  the  group  NH4 .  A  study 
of  this  reaction  in  the  laboratory  reveals  that  exactly  the  proportion  by 
weight  of  Ca3  (PO4)2  and  (NH4)3PO4,  expressed  in  this  equation  are,  in 
reality,  produced  by  nature. 

*  In  reality  a  mixture  of  secondary  and  tertiary  phosphate  of  calcium  is 
at  first  precipitated :  — 

3  Na2  HPO4  +  4  Ca  C12  =  Ca  HPO4  +  Ca3  (PO4)2  +  2  HC1  +  6  Na  Cl, 
and  in  consequence  of  the  free  hydrochloric  acid,  the  fluid  above  the  precipi- 
tate assumes  an  acid  reaction.      The  hydrochloric  acid,  however,  gradually 
reacts  with  the  tertiary  phosphate  to  produce  the  primary:  — 

Ca3  (P04)2  +  2  HC1  =  2  Ca  HPO4  +  Ca  Cl.,; 

so  that,  in  the  end,  the  reaction  given  above  is  realized.  The  same  is  true  of 
barium  phosphate. 

t   i.e.,  they  turn  red  litmus  of  a  blue  color  (page  75). 


METAPHOSPHATES  ;    PYKOPHOSPHATES.  231 

in  other  cases  the  tertiary  salt  alone  exists;  therefore  when,  for 
example,  silver  nitrate  is  added  to  the  secondary  phosphate  of 
sodium,  the  insoluble  tertiary  phosphate  of  silver  is  precipitated :  — 

Na2  HP04  +  3  Ag  N03  =  Ag3  P04  +  2  Na  N03  +  HN03 . 

This  and  similar  reactions  puzzled  the  chemists  of  former  days  not 
a  little ;  for,  by  mixing  a  solution  of  two  salts  which  were  neutral, 
they  obtained  a  liquid  of  acid  reaction  (because  of  the  free  acid 
formed). 

All  metaphosphates,  provided  they  are  not  those  of  a  volatile 
metal-like  substance,  such  as  ammonium,  are  unchanged  even  by 
quite  a  high  heat ;  it  follows  that  these  salts  will  be  formed  under 
conditions  which  render  the  existence  of  the  salts  of  the  great 
majority  of  other  acids  impossible;  as  a  consequence,  phosphoric 
acid  will,  if  the  temperature  is  sufficiently  increased,  ultimately 
decompose  the  salts  of  much  stronger  acids  (such  as  sulphuric), 
leaving  a  phosphate  in  their  place ;  while,  on  the  other  hand,  in 
solution,  or  in  the  cold,  the  exact  reverse  takes  place,  i.e.,  the  other 
acids  decompose  the  phosphates. 

Pyrophosphoric  acid  is  quadribasic,  but  forms  only  two  classes 
of  salts ;  namely,  those  with  M4  P2  07  and  those  with  M2  H2  P2  07  as 
their  general  formulae.  On  being  heated  with  water  the  pyrophos- 
phates  change  into  the  secondary  orthophosphates. 

Phosphoric  acid  is  necessary  for  animal  and  vegetable  life ; 
tertiary  calcium  phosphate  forms  the  major  portion  of  the  inorganic 
constituents  of  the  bones  and  teeth ;  but  phosphoric  acid,  combined 
in  some  form,  is  also  found  in  the  blood,  in  the  muscle  and  nerve 
tisues,  and  in  the  brain.  The  phosphates  which  are  found  in  the 
soil  are  generally  of  an  insoluble  variety ;  however,  the  chemical 
action  of  the  various  substances  which  are  present,  aided  by  water, 
renders  them  finally  partially  soluble  and  absorbable  by  plants. 
The  primary  calcium  phosphate  has,  when  mixed  with  other  ingre- 
dients, an  extensive  sale  as  superphosphate ;  it  is  used  as  a  fertilizer. 

Two  other  oxides  of  phosphorus,  P02  and  P40,  are  known; 
they  are  of  little  importance  except  that  P02  is  analogous  to  N02 . 
A  larger  work  must  be  consulted  for  their  description. 

One  more  acid  of  phosphorus,  hypophosphorous  acid,  H3P02, 
remains  to  be  considered.  Salts  of  this  acid  are  produced  by  the 
action  of  phosphine  on  a  solution  of  an  alkaline  hydroxide.  The 


232  HYPOPHOSPHOROUS    ACID. 

free  acid  is  formed  by  acidifying  the  barium  salt  with  the  calcu- 
lated quantity  of  diluted  sulphuric  acid.  It  is  a  crystalline  sub- 
stance which  melts  at  17°.  4,  and  which  oxidizes  to  phosphorous  acid 
when  in  contact  with  the  air.  When  heated  it  changes  to  phosphine 
and  phosphoric  acid. 

Hypophosphorous  acid,  like  phosphorous  acid,  is  a  hydrated  acid, 
which  is  derived  from  the  hypothetical  meta  acid  by  the  addition 
of  one  molecule  of  water  to  each  formula  weight  :  — 


Only  one  atom  of  hydrogen  in  a  molecule  of  hypophosphorous 
acid  is  capable  of  being  replaced  by  metals  to  form  a  salt  ;  this 
fact  is  exactly  what  would  be  expected  from  the  behavior  of 
phosphorous  acid.  (See  page  224.) 

The  following  table  shows  the  connection  between  the  compounds 
discussed  in  the  last  two  chapters  :  — 


OXY-ACIDS   OF   PHOSPHOKUS  ;    TABLE   OF. 


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234  .  ARSENIC  ;    OCCURRENCE. 


CHAPTER   XXX. 

ARSENIC   AND   ARSINE. 

Arsenic  ;  symbol,  As  ;  atomic  iv  eight,  75  ;  specific  gravity,  5.7.  Spe- 
cific gravity  of  vapor  at  red  heat,  air  =  1,  is  10.3,  H2  =  2,  is 
'296.6  ;  molecular  weight  of  As4  =  300,  of  As2  =  150.  Arsine, 
As  H3 ;  specific  gravity,  air  =  1,  is  2.7,  H2  =  2,  is  78.02 ; 
'molecular  weight,  78.021 ;  1  c.c.  of  the  gas  at  0°  and  .76  m. 
pressure  weighs  .003499  gram. 

ARSENIC  occurs  quite  frequently  in  the  form  of  the  uncombined 
element,  especially  in  formations  which  contain  the  metallic  sul- 
phide ;  the  native  arsenic  is  found  in  grayish  black,  reniform  masses. 
Combined  arsenic  occurs  in  the  arsenides  of  many  metals,  the  chief 
examples  of  which  are  :  — 

Arsenopyrite,  FeAsS  (corresponding  to  FeS2,  iron  pyrites,  one  atom  of 
sulphur  being  replaced  by  one  of  arsenic). 

Smaltite,  (Co,  Fe  Ni)  As2  (corresponding  to  Fe  S2 ,  iron  pyrites,  both  atoms 
of  sulphur  being  replaced  by  arsenic). 

Cobaltglance,  Co  As  S. 

Two  sulphides  of  arsenic  are   also  not  infrequently  found  as 

minerals ;  they  are  :  — 

As2S2  (realgar). 
As2S3  (orpiment). 

In  addition  to  these  occurrences,  arsenic  also  appears  as  the 
oxide  As203,  and,  in  some  arsenates,  these  latter  probably  the 
result  of  the  oxidation  of  the  arsenides.  The  arsenic  of  commerce 
is  either  the  naturally  occurring  element,  or  it  is  prepared  from  the 
arsenides  which  are  found  as  minerals. 

The  two  sulphides  referred  to  above  and  the  oxide  As2  03  have 
been  known  since  ancient  times ;  Aristotle  mentions  the  former,  but 
the  term  arsenicon  seems  first  to  have  been  used  by  Dioscorides. 
The  element  arsenic  does  not  seem  to  have  been  isolated  until  the 
latter  part  of  the  seventeenth  century;  of  course,  its  chemical 
nature  was  not  understood  until  much  later. 


ARSENIC  ;    PREPARATION,  PROPERTIES.  235 

The  commercial  preparation  of  arsenic  depends  upon  the  forma- 
tion of  ferrous  sulphide  and  free  arsenic,  when  arsenopyrite  is 
heated  :  -  Fe  As  S  =  Fe  S  +  As. 

The  impure  arsenic  so  prepared  is  purified  by  sublimation ;  for  when 
heated,  it  volatilizes  without  previously  melting,  and  then  collects 
in  crystals  on  the  colder  parts  of  the  retorts. 

Arsenic  has  a  steel-gray  color  and  metallic  lustre;  when  frac- 
tured, it  resembles  white  pig  iron.  Arsenic,  like  sulphur  and  phos- 
phorus, exists  in  two  allotropic  forms,  one  of  which  is  crystalline,* 
and  the  other  amorphous.  When  the  element  is  heated  in  a  glass 
tube,  amorphous  arsenic  is  deposited  on  the  walls  near  the  flame  as 
a  black  mirror,  while  the  bright,  shiny  crystals  of  the  other  variety 
appear  upon  the  cooler  portions.  The  amorphous  form  is  changed 
into  the  crystalline  one  by  heating  to  360°.  Crystallized  arsenic 
has  a  specific  gravity  of  5.76,  amorphous  of  4.71. 

When  heated  in  the  air,  arsenic  burns  to  form  the  trioxide 
As2  03 ;  in  this  particular  the  element  differs  from  phosphorus,  which, 
under  similar  circumstances,  forms  the  pentoxide  P2  05 .  When 
heated  or  burned,  arsenic  assumes  a  peculiar  odor,  somewhat  resem- 
bling that  of  phosphine.  The  element  unites  directly  with  chlorine, 
bromine,  or  iodine  to  form  the  corresponding  halogen  compounds, 
and  arsenic  also  readily  combines  with  a  number  of  metals,  thus 
yielding  arsenides. 

Arsenic  volatilizes  at  450° ;  the  vapors  have  a  lemon  yellow 
color  and  a  most  disagreeable  garlic  odor.  Their  specific  gravity, 
air  =  1,  at  a  low  red  heat,  is  10.3,  which,  H2  =  2,  gives  296.6  •  the 
molecular  weight  at  this  temperature  is  therefore  300  and  the  mole- 
cule consists  of  four  atoms  (As4).  The  density  of  the  arsenic 
vapors,  however,  gradually  diminishes  as  the  heat  is  increased, 
until  it  reaches  5.37  at  1736°. f  The  molecules,  As4,  which  corre- 
spond to  those  of  phosphorus,  P4 ,  therefore  begin  to  dissociate  as  a 
white  heat  is  reached,  so  that  they  very  nearly  change  into  As2 ;  $ 
—  the  value  for  the  specific  gravity  would  be  5.22,  air  =  1,  were 
complete  dissociation  into  As2  to  result. 

*  Hexagonal  system,  rhombohedra. 

t  Latest  determinations  by  H.  Biltz  and  Y.  Meyer.  (Ber.  d.  Deutsch. 
Chem.  Gesell.;  22,  726). 

t  Possibly  they  partially  dissociate  into  the  individual  atoms;  there  is  no 
means  of  determining  this  point. 


236  AESINE;    PREPARATION. 

Arsenic  is  an  element  which  is  on  the  boundary  line  between 
metal  and  not-metal.  Crystallized  arsenic  is  entirely  metallic  in 
appearance ;  although  it  is  neither  malleable  nor  ductile,  it  conducts 
heat  and  electricity  quite  readily.  Chemically,  arsenic  is  almost 
entirely  a  not-metal,  but  its  approach  to  a  metallic  character  is 
evinced  by  the  instability  of  its  hydrogen  compound. 

Arsine,  AsH3,  is  the  analogon  of  phosphine,  PH3,  and  of  am- 
monia, NH3 .  It  is  produced,  similarly  to  phosphine,  by  the  action 
of  an  acid  upon  some  arsenide,  thus :  — 

As2Zn3  +  6H  Cl  =  2  AsH3  +3  ZnCl2. 

Zinc  arsenide.  Arsine. 

Such  a  reaction  is  exactly  like  the  ones  which  were  studied  when 
the  preparation  of  hydrogen  sulphide  and  other  hydrogen  compounds 
was  discussed,  thus :  — 

Fe  S       +  2  H  Cl  =  Fe  C12     +  H2  S. 

Zn3  As2  +  6  H  Cl  =  3  Zn  C12  +  2  As  H3 . 

This  method  furnishes  pure  arsine.  Another  and  more  important 
way  of  preparing  arsine  is  by  the  action  of  nascent  hydrogen  upon 
an  acid  solution  of  a  soluble  arsenic  compound.  For  instance,  a 
solution  of  arsenic  trioxide  in  hydrochloric  acid,  when  added  to  zinc 
which  is  covered  with  dilute  sulphuric  acid  and  which  is  therefore 
generating  hydrogen,  will  develop  arsine,  the  latter,  however, 
naturally  is  mixed  with  hydrogen :  — 

As,03  +12H  =  2AsH3  +  3  H2  0. 

The  most  delicate  test  for  arsenic  (Marsh's  test)  is  based  upon  this 
chemical  fact.63 

Arsine  is  a  colorless  gas,  with  a  most  disagreeable  odor.  It  is 
an  intense  poison,  so  that  it  is  imperatively  necessary  to  take  every 
precaution  when  experimenting  with  the  gas,  especially  when  the 
latter  is  obtained  pure,  as  it  is  from  zinc  arsenide.*  The  specific 
gravity  of  arsine  is  2.7,  air  ==  1,  or  78.02,  H2=  2;  the  molecular 
weight  is,  therefore,  78.024  (As  =  75,  3  H  =  3.024 ;  As  H3  =  78.024) 
and  the  formula  AsH3.  Arsine  changes  to  a  liquid  at — 102°, 
and  forms  a  white  crystalline  mass  at  — 119°.  When  ignited  in 
the  air  the  gas  burns  with  a  bluish  white  flame,  forming  water  and 
arsenic  trioxide,  the  latter  compound  appearing  in  the  form  of  a 

*  Arsine  should  never  be  generated  otherwise  than  under  a  hood  or  in  the 
open  air. 


ARSINE;  PROPERTIES.  237 

white  smoke.  If  the  air  supply  in  which  arsine  is  burning  is 
limited,  or  if  the  flame  is  cooled,  for  example,  by  holding  a  porce- 
lain plate  in  it,  the  arsenic  formed  by  the  decomposition  of  the  gas 
into  its  elements  will  separate  and  will  form  a  dark  spot  on  the 
surface  in  contact  with  it.  A  mixture  of  arsine  and  oxygen  explodes 
violently  when  ignited. 

Arsine  is  much  less  stable  than  phosphine.  When  it  is  passed 
through  a  heated  glass  tube  it  readily  decomposes  into  hydrogen 
and  arsenic,  the  latter  being  deposited  as  a  black  mirror.  Of 
course,  arsine  is  a  powerful  reducing  agent ;  sulphuric  acid  is  readily 
decomposed  by  it,  just  as  that  acid  is  by  hydroiodic  acid  or  hydro- 
gen sulphide,  but  in  this  case  the  acid  is  robbed  of  all  of  its  oxygen 
while  the  sulphide  of  arsenic  is  formed  ;  *  other  acids  and  even  water 
or  an  alkaline  solution  can  likewise  decompose  hydrogen  arsenide. 
Arsine  reduces  silver  nitrate  in  solution ; .  metallic  silver  is  precipi- 
tated and  arsenic  trioxide  is  formed ;  this  fact  is  of  importance  in 
the  detection  of  arsenic.  The  basic  properties  manifested  by  ammo- 
nia and  phosphine  are  absent  in  arsine;  the  latter  forms  no 
arsonium  compounds ;  it  can  be  made  to  act  as  a  base  only  when 
the  hydrogen  atoms  contained  in  it  are  substituted  by  some  more 
positive  (so-called  organic)  radicle,  like  methyl,  f 

*  This  would  appear  to  be  an  example  of  the  nascent  action  of  a  solid 
element;  arsenic,  at  the  moment  of  its  liberation  from  arsine,  readily  com- 
bines with  the  hydrogen  sulphide,  or,  perhaps,  with  the  sulphur  formed  by 
the  reduction  of  the  sulphuric  acid. 

t  See  methane,  chapter  on  compounds  of  carbon  and  hydrogen. 


238  ARSENIC  ;    HALIDES    OF. 


CHAPTER   XXXI. 

THE  COMPOUNDS  OF  ARSENIC  WITH  THE  HALOGENS,  WITH 
OXYGEN,  AND  WITH  OXYGEN  AND  HYDROGEN. 

THE  halogen  compounds  of  arsenic  are  not  complicated  by  the 
existence  of  two  series,  for  only  the  trihalogen  derivatives  have 
with  certainty  been  prepared ;  their  chief  characteristics  are  as 
follows :  — 

As  F3 ,  arsenic  trifluoride,  liquid,  boils  at  63°.* 
AsCl3,  arsenic  trichloride,  liquid,  boils  at  130°.2,  solid  at  -18°. 
'AsBr3,  arsenic  tribromide,  solid,  melts  at  20-25°,  boils  at  220°. 
As  I3 ,  arsenic  tri-iodide,  solid,  sublimes  when  heated. 

Arsenic  shows  its  resemblance  to  the  metals  and  at  the  same 
time  its  connection  with  the  not-metals  nowhere  better  than  in  the 
chemical  behavior  of  its  chlorine  compound.  The  latter  substance 
can  be  formed,  as  are  the  chlorides  of  metals,  by  the  action  of  hydro- 
chloric acid  on  the  trioxide  of  arsenic,  so  that  in  this  case  arsenic 
trioxide  is  a  base  and  arsenic  a  metal :  - 

As203  +  6HCl  =  2  AsCl3  +  3H20. 

This  formation  of  the  chloride  of  arsenic  is  not,  however, 
possible  unless  very  little  water  is  present ;  it  takes  place,  for  in- 
stance, when  dry  hydrochloric  acid  is  passed  over  the  oxide  of 
arsenic,  or  when  the  latter  substance  is  distilled  with  a  mixture  of 
sulphuric  and  hydrochloric  acid.t  On  the  other  hand,  when  an 
excess  of  water  is  added  to  the  trichloride,  it  is  entirely  decomposed 
into  hydrochloric  acid  and  arsenic  trioxide :  — 

*  A  pentafluoride  of  arsenic  (AsF5)  has  been  isolated  by  Moissan. 
This  compound  was  prepared  by  electrolyzing  the  trifluoride,  when  arsenic 
and  the  pentafluoride  are  formed.  The  compound  is  interesting  because  it  is 
the  only  pentahalide  of  arsenic,  and  therefore  brings  the  halides  of  arsenic  in 
line  with  those  of  phosphorus. 

t  The  trichloride  of  arsenic  is  also  formed  when  the  trioxide  is  warmed 
with  a  concentrated  solution  of  hydrochloric  acid  ;  it  then  separates  as  an  oil, 
insoluble  in  the  excess  of  the  acid. 


ARSENIC   TRIOXIDE.  239 


a.  2AsCl3  +  6H20=2As(OH)3  + 

b.  2As(OH)3=As203-j-3H20.* 

An  arsenious  acid  which  corresponds  to  phosphorous  acid  does 
not  exist;  where  its  formation  is  to  be  expected,  not  it,  but  its 
anhydride,  As203,  results.  The  easy  decomposition  of  the  trichlo- 
ride of  arsenic  by  an  excess  of  water,  therefore,  is  a  phenomenon 
classing  arsenic  as  a  not-metal,  for  the  chlorides  of  the  pronounced 
metals  are  unchanged  by  the  action  of  water.  f  The  addition  of  a 
little  water  only  partially  decomposes  arsenic  trichloride  :  — 

rei  +  HOH  (OH 

As  •]  Cl  +  HOH     =     As  ]  OH  +  2  HC1. 

(  ci  I  ci 

The  compound  so  produced  is  a  so-called  basic  salt,  i.e.,  a  salt  which 
is  in  part  hydroxide  and  in  part  chloride  (see  antimony  trichlo- 
ride). The  chloride,  bromide,  or  iodide  of  arsenic  can  readily  be 
formed  by  the  direct  union  of  the  elements.  These  compounds  are 
all  extremely  poisonous  substances  ;  the  fluoride,  chloride,  and 
bromide  fume  in  the  air,  absorb  moisture,  and  decompose,  leaving 
the  trioxide. 

Arsenic  forms  two  oxides,  As203,$  and  As2O5;  they  are  the 
anhydrides  respectively  of  arsenious  and  of  arsenic  acid  ;  arsenious 
acid  is,  however,  known  only  in  its  salts,  for  we  have  seen  that  when 
it  is  liberated  from  these  it  at  once  breaks  down  into  its  anhydride 
and  water.  Owing  to  the  few  reactions  in  which  arsenic  acts  like 
a  metal,  the  oxides  are  sometimes  named  arsenious  and  arsenic 
oxides  in  conformity  with  the  nomenclature  usually  adopted  where 
a  metal  forms  two  such  compounds  (see  page  26). 

Arsenious  oxide  (arsenic  trioxide,  As203)  is  the  most  common 
preparation  of  arsenic,  having  been  known  to  the  ancients,  and  hav- 
ing been  a  familiar  substance  ever  since  the  time  of  the  Romans  ; 
it  is  popularly  known  by  the  name  of  arsenic  or  white  arsenic  ; 
when  cases  of  poisoning  by  arsenic  occur,  the  substance  used  is 

*  An  intermediary  product,  As(OH)2  Cl,  is  formed,  in  all  probability 
before  complete  decomposition  is  accomplished.  This  compound  is  a  basic 
chloride  of  arsenic  (see  page  181).  t  See  page  79. 

t  The  specific  gravity  of  this  substance,  taken  by  Victor  Meyer,  above  1,500°, 
corresponds  to  the  molecule  of  the  formula  As4  O6  .  The  smallest  molecule  of 
the  trioxide  of  arsenic  with  which  we  are  acquainted  is  therefore  As4  O6  . 


240  ARSENIC   TKIOXIDE;   ACTION   AS   A   POISON. 

generally  "  white  arsenic."  As  its  name  implies,  it  is  a  white  solid, 
resembling  ordinary  flour.  The  commercial  product  is  formed  by 
roasting  the  arsenical  sulphides  which  occur  as  mineral  deposits ; 
arsenico-pyrite  is  especially  advantageous  for  this  purpose ;  in 
roasting  this,  care  must  be  taken  to  have  a  sufficient  supply  of  air, 
otherwise  arsenic  (formed  as  shown  on  page  235)  sublimes  together 
with  the  oxide.  Quantities  of  white  arsenic  are  also  produced 
while  burning  cobalt,  nickel,  tin,  and  silver  ores ;  the  arsenious 
oxide  is  collected  in  cold  chambers  and  is  purified  by  sublimation. 
The  white  arsenic  of  commerce  generally  contains  traces  of  the 
corresponding  oxide  of  antimony. 

Sublimed  arsenic  trioxide  appears  in  two  forms,  in  the  first  as  a 
dimorphous  substance,  crystallizing  both  in  the  regular  and  in  the 
monoclinic  system,  and  in  the  second  as  an  amorphous,  glass-like 
body,  which  gradually  changes  into  a  porcelain-like  mass.  The 
oxide  volatilizes  at '200°.* 

Arsenious  oxide  is  a  poison  to  all  animals,  and  even  to  plants. 
Owing  to  its  t^stelessness  and  the  readiness  with  which  it  can  be 
mixed  with  foods,  it  is  frequently  used  as  a  poison,  both  intention- 
ally and  accidentally.  It  acts  in  two  ways. 

a.  As  a  corrosive  substance,  attacking  the  organic  surfaces  with 
which  it  comes  in  contact,  it  therefore  causes  local  inflammation  in 
the  stomach  and  intestinal  tract. 

b.  It  has  a  destructive  effect  on  the  medullas  of   the  nerves. 
The  more  rapidly  the  poison  is  absorbed  the  less  the  action  under  a 
and  the  more  the  one  under  b  is  observed ;  it  follows  that  the  latter 
is  more  prominent  when  the  arsenic  is  administered  in  solution, 
and,  of  course,  it  is  of  material  influence  whether  the  stomach  is  full 
or  empty  at  the  time  of  taking.     .005  gram  has  a  marked  effect, 
and  if  continued  can  cause  death;  .06  to  .12  gram  may  cause  death 
in  a  few  days,   and  .2   to  .3   gram  in  a  few  hours.     Symptoms: 
nausea,  salivation,  burning  in  the  gastric  region,  vomit  yellow  or 
greenish,  and  possibly  streaked  with  blood,  while  traces  of  white 
arsenic  may  be  visible  therein,  great  thirst,  colic,  and  sensitiveness 
of  the  abdomen.     The  symptoms  of  the  action  under  b  are  :  intense 
fear,  convulsive  movements,  trembling  and  cramps  in  the  extremi- 
ties, fainting  spells,  and  delirium ;  where  large  quantities  have  been 

*  According  to  Selmi,  it  already  begins  to  vaporize  at  100°. 


ARSENITES;    ARSENIC    PENTOXIDE.  241 

rapidly  absorbed  the  patient  may  be  entirely  unconscious.  On  post 
mortem  examination,  fatty  degeneration  of  the  liver  and  heart,  as 
well  as  in  many  other  organs,  the  effects  of  the  poisoning  being  much 
like  those  resulting  from  the  taking  of  phosphorus.  An  antidote  is 
a  mixture  of  freshly  precipitated  ferric  hydroxide  and  magnesium 
oxide,  the  endeavor  being  to  form  the  insoluble  ferric  arsenite. 
Glass-like  arsenic  trioxide  is  much  more  soluble  in  water  than  is  the 
crystallized  variety,  100  parts  of  water  at  ordinary  temperatures 
dissolve  four  parts  of  the  former  and  1.2  to  1.3  parts  of  the  latter. 

The  chemical  action  of  arsenic  trioxide  is  twofold,  for  it  can 
appear  both  as  a  base  and  as  the  anhydride  of  an  acid. 

a.  It  acts  like  a  base,  because  it  can  dissolve  in  a  number  of 
acids.  One  or  two  of  the  salts  produced,  for  instance,  the  chloride, 
AsCl3,  and  a  sulphate  (As2  (S04)  3-f-  S03),  have  been  isolated; 
these  substances  are  decomposed  by  water. 

ft.  It  acts  like  the  anhydride  of  an  acid,  because  it  dissolves  in 
bases  to  form  arsenites  :  — 

As,  03  +  2  KOH  =  2  As  02  K  +  H2  0, 

The  meta-arsenites  ( like  potassium  meta-arsenite,  As  02  K  )  are  the 
most  common  salts  of  this  acid;  they  correspond  to  the  nitrites 
N02  M ;  a  few  orthoarsenites,  As  03  M3 ,  are  also  known.  The 
arsenites  of  the  alkali  metals  are  soluble  in  water ;  the  others  are 
either  soluble  with  difficulty  or  entirely  insoluble ;  they  are  decom- 
posed by  hydrochloric  acid.  The  alkaline  solutions  of  arsenic 
trioxide  are  most  powerful  reducing  agents ;  they  have  the  greatest 
tendency  to  take  up  oxygen  to  form  arsenates,  but,  on  the  other 
hand,  arsenious  oxide  is  quite  readily  reduced,  and  a  method  for 
detecting  arsenic  is  based  upon  the  ease  with  which  this  substance 
gives  up  oxygen.64 

Arsenic  pentoxide,  As20R,  is  the  anhydride  of  arsenic  acid;  it 
corresponds  to  N2  05  and  P2  05 ;  three  acids,  analogous  to  those  of 
phosphorus,  are  derived  from  this  oxide ;  they  are :  — 

H  As  03 ,  meta-arsenic  acid, 
H3  As  04 ,  orthoarsenic  acid, 
H4  As2  07 ,  pyroarsenic  acid. 

Arsenic  pentoxide  is  produced  when  arsenic  acid  is  heated  to  a 
low  red  heat ;  a  higher  temperature  produces  decomposition  into 


242  ARSENIC   ACIDS. 

arsenic  trioxide  and  oxygen.      It  is  a  colorless,  amorphous  mass, 
which  greedily  absorbs  moisture  and  finally  deliquesces. 

Orthoarsenic  acid,  H3  As  04  ,  the  only  arsenic  acid  which  exists 
in  aqueous  solutions,  is  produced  by  oxidizing  arsenic  trioxide  (for 
instance,  by  chlorine,  bromine,  or  nitric  acid*).  When  the  solu- 
tion produced  by  oxidizing  arsenic  trioxide  is  evaporated  to  dryness, 
orthoarsenic  acid  separates  in  needle-like  crystals  ;  when  these  are 
heated  to  180°  they  separate  water,  and  change  into  pyroarsenic 


which  latter,  at  206°,  changes  into  meta-arsenic  acid  :  — 
H4As207  =  2HAs03  +  H20, 

and  this  substance  finally,  at  red  heat,  decomposes  to  yield  the  an- 
hydride, As2  05  .  Each  one  of  these  compounds  is  converted  into 
orthoarsenic  acid  by  solution  in  water. 

Orthoarsenic  acid  is  tribasic,  and  forms  primary,  secondary,  and 
tertiary  salts  ;  the  formulae  of  these  are,  of  course,  exactly  parallel 
to  those  of  the  corresponding  derivatives  of  phosphoric  acid  (see 
page  229).  Primary  and  secondary  arsenates,  when  they  are  heated 
to  redness,  undergo  the  same  changes  which  take  place  with  primary 
and  secondary  phosphates  :  — 

MH2  As  04  =  M  As  03     +  H2  0. 

Primary  arsenate.          Meta-arsenate. 

M2  H  As  04         =  M4  As2  0T     +  H2  0. 

Secondary  arsenate.       Pyroarsenate. 

The  arsenates  of  the  alkalies  are  soluble  in  water,  but  only 
the  primary  arsenates  of  the  other  metals  dissolve  ;  the  latter 
are,  however,  dissolved  by  mineral  acids  (see  page  229)  ;  it  will  be 
remembered  that  this  is  also  the  case  with  the  phosphates. 

Arsenic  acid  is  a  tolerably  good  oxidizer  ;  quite  a  number  of 
reducing  agents  reduce  it  to  arsenic  trioxide.  It  has  of  late  years 
been  extensively  used  in  the  manufacture  of  aniline  dyes,  because, 
while  it  certainly  gives  up  its  oxygen,  it  does  not  do  so  with  such 
facility  as  to  destroy  the  organic  substance  it  acts  upon.  Sulphur 
dioxide  reduces  arsenic  acid  as  follows  :  — 

*  The  pupil  should  write  these  equations,  using  the  knowledge  acquired 
in  studying  the  previous  oxidizing  action  of  these  substances. 


AKSENIC   ACIDS.  243 

H3  As  04  +  H2  0  +  S02  =  H3  As  03  +  H2  S04 , 

the  arsenious  acid  so  formed,  of  course,  breaking  down  into  arsenic 
trioxide  and  water.  This  reaction  serves  to  distinguish  arsenic  acid 
from  phosphoric  acid,  for  the  latter  substance,  although  in  every 
respect  like  the  former,  has  no  tendency  whatever  to  give  up  oxygen 
and  change  to  phosphorous  acid.  Arsenic  acid  is  a  powerful  poison. 


244  ARSENIC;   TEISULPHIDE. 


CHAPTER   XXXII. 

THE  COMPOUNDS  OF  ARSENIC  WITH  SULPHUR,  AND  WITH 
SULPHUR  AND  HYDROGEN. 

ARSENIC  forms  three  sulphides,  two  of  which  correspond  in 
formula  to  the  two  oxides  ;  they  are  As2  S3  and  As2  S5 ,  while  the 
third  is  a  ruby  red  mineral  substance  (realgar),  As2  S2 . 

Arsenic  trisulphide  and  arsenic  pentasulphide  are  especially 
interesting,  because  they  act  much  like  anhydrides  of  oxy-acids,  and 
therefore  illustrate  the  marked  chemical  resemblance  between  sul- 
phur and  oxygen. 

Arsenic  trisulphide  is  found  as  a  natural,  yellow  colored  mineral 
which  bears  the  name  of  orpiment.  It  is  produced,  as  are  the  sul- 
phides of  many  metals,  by  the  action  of  hydrogen  sulphide  upon 
an  acidified  solution  of  arsenic  trioxide,  for  the  sulphide  of  arsenic 
is  one  of  the  sulphides  which  are  insoluble  in  dilute  acids.  (See 
page  100.)  2  As  C13  +  3  H2  S  =  A2  S3  +  6  H  Cl. 

When  precipitated,  it  is  a  bright  yellow  powder.  Arsenic  trisul- 
phide can  also  be  formed  by  direct  union  of  the  elements,  just  as  is 
the  trioxide.  Arsenic  trioxide  dissolves  in  alkaline  solutions  to 
form  a  meta-arsenite  :  — 

As2  03  +  2  KOH  =  2  As  02  K  +  H2  0, 

and  the  trisulphide  is  dissolved  by  hydrosulphides  in  exactly  the 
same  way:-  As2S3 -f  2KSH  =  2  As  S2K  +  H2S  ; 
in  this  case  potassium  metasulpharsenite  is  formed.  Metasulphar- 
senious  acid  may  be  considered  as  derived  from  an  orthosulphar- 
senious  acid  by  the  separation  of  hydrogen  sulphide,  just  as 
meta-arsenious  acid  might  be  formed  from  orthoarsenious  acid. 
The  following  equations  will  make  this  clear :  — 

r  SH  (S 

As  •<  SH  =  As  •<  OTT  r  H2  S, 

(SH 
Orthoaulpharsenious  acid  =  Metasulpharsenious  acid  +  Hydrogen  sulphide ; 


ARSENIC  ;   PENTASULPHIDE.  245 

•H 


Orthoarsenious  acid  =  Meta-arsenious  acid  +  Water. 

We  saw,  however,  that  meta-arsenious  acid  is  incapable  of  exist- 
ence ;  for,  when  liberated  from  its  salts,  it  at  once  breaks  down  as 
follows  :  -  2  As  02  H  =  As2  03  +  H2  0  : 

metasulpharsenious  acid  acts  in  exactly  the  same  way  :  — 

2  AsS2H  =  As2S3  +  H2S; 

and,  therefore,  when  acids  are  added  to  solutions  of  the  sulpharsen- 
ites,  arsenic  trisulphide  is  precipitated  :  — 

2  AsS2K  +  2  HC1  =  As2S3  +  H2S+2  KC1. 

Arsenic  trisulphide  can  readily  be  dissolved  in  the  hydroxides  of 
the  alkalies,  for  then  a  mixture  of  the  arsenite  and  sulpharsenite  is 
produced  ;  *  it  is  also  dissolved  by  ammonium  sulphhydrate,  but  in 
the  latter  event  a  salt  of  pyrosulpharsenious  acid  is  the  result  ;  the 
free  acid  corresponding  to  this  ammonium  salt  does  no't  exist,  but  it 
can  be  supposed  to  be  formed  by  the  separation  of  one  molecule  of 
hydrogen  sulphide  from  two  formula  weights  of  an  hypothetical 
orthosulpharsenious  acid  :  — 

SH      HS  SH      HS 


(SH      HS)  (_._S--) 

2  As  (SH  )3  =  As2  S5  H4  +  H2  S.f     (  See  page  228.) 

Some  salts  derived  from  orthosulpharsenious  acid,  As  (  SH  )3  ,  are 
also  known  ;  none  of  the  acids  exist  in  the  free  state  ;  they  at  once 
break  down  into  the  trisulphide  of  arsenic  and  hydrogen  sulphide 
when  liberated  from  their  salts. 

Arsenic  pentasulphide  can  be  produced  from  the  trisulphide  by 
fusing  the  latter  with  sulphur  :  - 

As2  S3  -f  2  S  =  As2  S5  ; 

this  sulphurization  corresponding  to  the  oxidation  of  arsenic  tri- 
oxide:-  As203  +  2  0=As205. 

Arsenic  pentasulphide  cannot  be  precipitated  from  cold  solutions  of 
pure  arsenic  acid  by  means  of  hydrogen  sulphide,  because  the  ar- 

*  2  A&J  S3  +  4  KOH  =  3  As  S2  K  +  As  O.2  K  +  2  H2  O. 
t  As2  S3  +  4  NH4  SH  =  As2  S5  (  NH4  )4  +  2  H2  S. 


246  SULPHAK  SENATES. 

senic  acid  is  first  reduced  to  arsenious  acid  by  means  of  sulphuretted 
hydrogen,  which  latter  substance  is  oxidized  to  sulphur  ;  as  a  conse- 
quence, a  mixture  of  arsenic  trisulphide  and  sulphur  is  produced ;  * 
the  pentasulphide  is,  however,  produced  by  the  action  of  hydrogen 
sulphide  on  a  boiling  solution  of  arsenic  acid. 

Arsenic  pentasulphide  dissolves  in  the  sulphides,  sulphhydrates, 
or  hydroxides  of  the  alkalies  to  form  sulpharsenates  ;  it  therefore 
acts  as  the  anhydride  of  sulpharsenic  acid :  — 

As2  S5  +  2  K  SH  =2  As  S3  K  +  H2  S. 

The  sulpharsenates  are  derived  from  three  sulpharsenic  acids, 
orthosulpharsenic  acid,  As  S4  H3 ,  corresponding  to  orthoarsenic 
acid,  As  04  H3 ;  metasulpharsenic  acid,  As  S3  H,  corresponding  to 
meta-arsenic  acid,  As  03  H ;  and  pyrosulpharsenic  acid,  As2  S7  H4 , 
corresponding  to  pyroarsenic  acid,  As2  07  H4 .  When  an  acid  is 
added  to  a  solution  of  a  sulpharsenate,  orthosulpharsenic  acid  is 
precipitated :  — 

As  S4  K3  +  3  H  Cl  =  As  S4  H3  +  3  K  Cl. 

This  orthosulpharsenic  acid,  when  boiled,  changes  to  the  pentasul- 
phide of  arsenic  and  hydrogen  sulphide,  just  as  the  orthoarsenic 
acid  changes  to  arsenic  pentoxide  and  water  at  a  red  heat :  — 

2  As  S4  H3  =  As2  S5  +  3  H2  S, 
2  As04H3=As205-f  3H20. 

When  arsenic  trisulphide  is  dissolved  in  the  sulphides  of  the  alkalies 
which  contain  an  excess  of  dissolved  sulphur,!  then  the  sulpharsen- 
ates are  found  in  the  solution ;  the  superfluous  sulphur  sulphurizes 
the  trisulphide  just  as  oxygen  oxidizes  the  trioxide  ;  the  penta- 
sulphide so  formed,  of  course,  produces  the  sulpharsenate  by  union 
with  the  metallic  sulphide  :  — 

As2S5  +K2S=2  AsS3K. 

The  following  table  shows  the  most  important  facts  discussed  in 
this  chapter  :  — 

*  H3  As  O4  +  H2  S  =  H3  As  O3  +  H2  O  +  S.  If  the  solution  of  arsenic  acid 
is  strongly  acid  with  hydrochloric  acid,  then  arsenic  pentasulphide  is  precipi- 
tated by  means  of  hydrogen  sulphide  (Brauner  and  Tomicek;  Monatshefte 
fur  Chemie  ;  1887,  607).  In  any  event,  even  if  the  arsenic  acid  is  perfectly  pure, 
only  a  partial  reduction  takes  place.  See  also  McCay,  Am.  Chern.  Journ.  12,  547. 

t  See  foot-note,  page  155. 


ARSENIC  ;    TABLE   OF    OXIDES   AND   SULPHIDES. 


247 


OXIDES. 

As2  O3 ,  forms  no  acids,  forms  meta-arsenites,  M  As  O2 ,  and  orthoarsenites, 

M3AsO3. 

(  As  O3  H,  meta-arsenic  acid,  forms  meta-arsenates. 
As2  O5  <  As  O4  H3  ,  orthoarsenic  acid,  forms  orthoarsenates. 

(  As2  O7  H4 ,  pyroarsenic  acid,  forms  pyroarsenates. 

HALOGEN  COMPOUNDS. 

Arsenic  forms  trihalides,  As  X3,  but  no  pentahalides,  As  X5,  excepting  the  fluoride. 
The  trihalides  are  decomposed  by  an  excess  of  water,  forming  arsenic  trioxide 
and  halhydric  acids :  2  As  X3  +  3  H2  O  =  As2  O3  +  6  HX. 

SULPHIDES. 

As2  S3 ,  forms  no  acids,  forms  metasulpharsenites,  M  As  S2  ,  orthosulpharsen- 
ites,  M3  As  S3 ,  and  pyrosulpharsenites,  M4  As2  S5 . 

As,S5,  forms  orthosulpharsenic  acid,    ^  M  As  S3  ,  metasulpharsenates. 
As  S4  Hg ,  which  yields :  )  ^ As  S* <  orthosulpharsenates. 

J  (  M4  As2  S7 ,  pyrosulpharsenates. 


AESENITE8  AND  SULPIIAKSEMITES. 

AB8ENATE8  AND  8ULPHAE8ENATE8. 

Meta-salts. 
Ortho-salts. 
Pyro-salts. 

M  As  O2  ,  M  As4  S2. 
M3  As  O3  ,  M3  As  S3  . 

AT     \  "    S 

M  As  O3  ,  M  As  S3  . 

M3  As  O4  ,  M3  As  S4  . 
M4  As2  O7  ,  M4  As2  S7  . 

,   JJ.4  AB]  05  . 

248  ANTIMONY;  OCCURRENCE,  PREPARATION. 


CHAPTER  XXXIII. 

ANTIMONY   AND   STIBINE.      THE   COMPOUNDS   OP   ANTI- 
MONY  WITH   THE   HALOGENS. 

Antimony  ;  symbol,  Sb ;  atomic  weight,  120 ;  specific  gravity,  6.7. 
Specific  gravity  of  vapor,  at  1640°,  air  =  1,  is  9.78  ;  H2  =  2  is 
281.66  ;  molecular  iveight  of  Sb2  =  240.  Stibine ;  formula, 
Sb  H3 ;  specific  gravity  not  determined. 

BUT  little  antimony  is  found  as  the  native  *  element ;  it  is  most 
frequently  encountered,  combined  with  sulphur,  in  the  mineral  stib- 
nite,  Sb2  S3 ,  a  lead  gray  substance  with  metallic  lustre,  from  which 
most  of  the  antimony  of  commerce  is  obtained.  Stibnite  has  been 
known  since  the  most  ancient  times.  It  is  mentioned  by  Diosco- 
rides  as  trrt/x/xt  and  by  Pliny  as  stibium.  It  was  mainly  used  in 
medicine  as  an  external  application,  but  it  also  formed  a  pigment  for 
blackening  the  eyebrows.  The  name  antimonium  was  applied  to  it 
at  a  later  date.  The  element  and  its  compounds  interested  the 
immediate  successors  of  the  alchemists  greatly,  for  they  thought 
them  to  be  most  wonderful  and  potent  medicinal  remedies. 

Antimony  is  prepared  from  its  sulphide  by  one  of  two  common 
metallurgical  processes ;  the  compound  is  either  melted  with  iron, 
by  which  means  ferrous  sulphide  and  antimony  are  produced ;  or  it 
is  roasted  in  a  draught  of  air,  the  sulphur  burned  off,  and  the  re- 
sulting oxide  of  antimony  ( Sb2  04 )  further  heated  with  charcoal, 
whereupon  antimony  and  carbon  monoxide  are  formed. 

Antimony  is  much  more  metallic  in  its  nature  than  is  arsenic ; 
its  appearance  indicates  this,  for  it  is  silver  white,  with  a  metallic 
lustre  ;  the  metal  is  neither  malleable  nor  ductile,  has  a  crystalline 
structure,  f  is  brittle  and  easily  pounded  into  a  fine  powder ;  it 
melts  at  425°  and  is  vaporized  at  a  high  red  heat,  t  The  specific 

*  A  "  native"  element  is  an  element  occurring  uncombined  as  a  mineral, 
t  An  amorphous  form  of  antimony  has  also  been  described. 
J  This  vaporization  is  much  retarded  if  the  element  is  covered  with  a 
layer  of  oxide. 


ANTIMONY;  ALLOYS.  249 

gravity  of  vapor  of  antimony  at  1640°  is  9.78.  *  This  number  is 
somewhat  greater  than  the  one  which  should  be  found  were  the 
molecule  of  antimony  Sb2 ,  so  that  antimony  certainly  has  no  stable 
molecules  of  the  formula  Sb4 ,  corresponding  to  those  of  arsenic  and 
phosphorus,  As4  and  P4.  Probably,  were  it  possible  to  ascertain 
the  specific  gravity  of  antimony  at  about  1800°,  we  would  find  a 
value  which  would  indicate  molecules  formed  of  Sb2,  or  perhaps 
even  of  the  individual  atoms. 

Antimony,  when  heated  to  a  high  red  heat  in  air  or  in  oxygen, 
burns  to  form  the  trioxide,  Sb2  03;  it  unites  with  the  halogens  in 
the  same  way,  powdered  antimony  even  burns  vigorously  when 
dropped  into  a  flask  containing  chlorine,  while  the  trichloride, 
SbCl3,  is  produced.  Antimony  shows  its  metallic  nature,  chemi- 
cally, by  dissolving  in  hydrochloric  or  sulphuric  acid;  with  the 
former,  it  produces  the  trichloride  and  hydrogen,  with  the  latter  the 
sulphate  of  antimony  and  sulphur  dioxide,  for,  as  dilute  sulphuric 
acid  does  not  attack  the  element,  reduction  of  the  hot  and  concen- 
trated acid  takes  place  exactly  as  it  does  when  that  substance  is 
heated  in  contact  with  copper.  (See  pages  75,  137  and  151.) 

Nitric  acid  oxidizes  antimony,  as  it  does  phosphorus  or  arsenic, 
but  the  reaction  varies  according  to  the  temperature,  concentration, 
and  mass  of  the  acid ;  where  the  latter  is  cold  and/dilute,  antimony 
trioxide  is  formed ;  as  the  temperature  is  increased,  or  the  acid 
becomes  m^re  concentrated,  antimonic  acid  begins  to  be  produced 
until,  under  proper  conditions,  this  may  be  the  entire  result  of  the 
oxidation ;  of  course,  the  nitric  acid  is  reduced  at  the  same  time. 
(  See  pages  199  and  206.) 

Antimony  is  an  ingredient  of  a  number  of  commercially  very 
important  alloys. 

A  combination  produced  by  fusing  two  or  more  metals  together 
is  termed  an  alloy.  Some  meWla  can  be  alloyeH.  with  each  other  in 
ail  proportions,  others  only  partially,  while  sdme  wiM  not  mix  at 
all ;  an  exactly  parallel  case  is*  found  in  the  behavior  of  ordinary 
liquids,  some  of  which,  like  alcohol  .and  water,  can  foe  mixed  in  any 
ratio,  others,  like  water  and  ether,  will  only  partially  dissolve  each 
other,  while  lastly,  some  oils  and  water  remain"  entirely  separate. 
The  question  as  to  whether  alloys  are  mere*  mechanical  mixtures  of 

*  H.  Biltz  and  Y.  Meyer;  Ber.  d.  Deutsch.  Chem.    Gesell.;  22,  726. 


250  ALLOYS;   CHEMICAL   NATUKE   OF. 

fused  metals,  or  whether  they  have  the  character  of  chemical  com- 
pounds, has  been  the  subject  of  continued  discussion,  and,  indeed, 
the  same  may  be  said  of  solutions,  which  latter  certainly  have  not 
the  characteristics  of  mere  mechanical  mixtures.  The  facts  sustain- 
ing the  theory  of  chemical  combination  are  as  follows.  An  alloy 
has  a  specific  gravity  different  from  the  mean  of  the  specific  gravi- 
ties of  the  component  metals  and  a  melting  point  which  also  is  not 
the  mean  of  the  melting  points  of  constituents.  The  melting  point 
of  a  metal  is  generally  diminished  by  being  alloyed ;  some  alloys 
become  liquid  at  a  temperature  much  below  that  at  which  any  of 
the  constituents  fuse.  Heat  is  given  off  in  the  formation  of  alloys,* 
while  many  also  have  a  definite  crystalline  form.  ^Nevertheless,  those 
alloys  which  are  produced  by  metals  which  mix  in  any  proportion, 
do  not  have  that  characteristic  which  is  supposed  to  be  the  sine  qua 
non  of  a  chemical  compound  —  the  definite  composition  by  weight  — 
and  yet  they  cannot  be  separated  into  their  constituent  parts  by  sim- 
ple mechanical  means.  These  discrepancies  are  explained  by  the 
theory  that  certain  definite  compounds  of  the  metals  are  really 
formed,  but  that  these  are  further  dissolved  in  the  excess  of  one  or 
the  other  of  the  constituents.  This  theory  is  borne  out  by  the  fact 
that  many  alloys,  on  being  slowly  cooled,  allow  mixtures  of  a  defi- 
nite crystalline  form  and  gravimetric  composition,  to  separate  in  the 
same  way  as  ordinary  solid  chemical  compounds  can  be  ciystallized 
from  solutions. f  Those  metals  which  are  chemically  most  like 
each  other,  generally  mix  most  easily  to  form  alloys.  The  majority 
of  metals  are  white  or  gray,  and  most  alloys  are  also  white  or  gray ; 
copper  and  gold  are  red  and  yellow  respectively ;  their  alloys  pre- 
sent modifications  of  these  colors,  unless  the  admixture  of  the 
a,dded  metal  is  so  great  in  quantity  as  to  entirely  conceal  this  char- 
acteristic. $  Amalgams,  as  we  have  already  seen,  are  alloys  con- 
taining mercury.  They  are  easily  fused  if,  indeed,  they  are  not 
.soft  at  ordinary  temperatures ;  the  mercury  evaporates  when  the 
amalgam  is  heated  to  the  boiling  point  of  that  element.  §  Some 

*  Fused  zinc  and  fused  copper  when  mixed  with  each  other  liberate  so 
much  heat  that  the  mass  may  spatter  out  of  the  containing  crucible,  and  a 
large  part  of  the  zinc  evaporates.  It  must  also  be  remembered  that  heat  is 
not  infrequently  given  off  in  the  formation  of  mere  solutions. 

t  Rudberg;  Poggendorff  s  Annalen  18,  page  240. 

t  An  admixtureof  30%  of  tin,  to  copper,  will  entirely  destroy  the  red  color. 

§  In  that  way  amalgams  resemble  crystals  with  water  of  crystallization. 


STIBINE  ;    PREPARATION,  PROPERTIES.  251 

amalgams  have  a  definite  crystalline  form  and  chemical  composition. 
If,  as  seems  unavoidable,  we  regard  alloys  as  chemical  compounds, 
we  must,  nevertheless,  believe  that  they  certainly  cannot  be  classed 
with  substances  ordinarily  considered  as  such ;  011  the  other  hand, 
their  existence  is  a  constant  argument  against  a  too  dogmatic  con- 
ception of  the  laws  of  definite  proportions  and  of  valence. 

The  most  important  alloy  of  antimony  is  composed  of  one  part 
of  that  metal  to  four  of  lead ;  this  substance  is  used  as  type  metal, 
copper  and  bismuth  are  occasionally  added  to  this ;  the  metal  used 
for  stereotyping  also  contains  tin. 

The  hydrogen  compound  of  antimony  is  called  stibine.  It  is 
produced,  like  arsine,  by  the  action  of  acids  on  the  alloy  of  anti- 
mony and  zinc,*  of  by  the  action  of  nascent  hydrogen  on  soluble 
compounds  of  antimony ;  the  gas  is  therefore  produced  by  the  same 
means  which  furnish  arsine,  as  a  consequence  the  latter  may  con- 
tain stibine  unless  care  is  taken  to  exclude  that  gas ;  in  fact,  stibine 
may  be  mistaken  for  arsine  unless  especial  precautions  are  taken  to 
distinguish  between  the  two.65 

Stibine  is  a  colorless  and  odorless  gas,  scarcely  soluble  in  water. 
As  would  be  expected,  owing  to  the  metal-like  nature  of  antimony, 
this  gas  is  much  less  stable  than  arsine ;  it  can  be  compared  to 
hydrogen  telluride  (page  104)  in  this  respect,  for  it  partially  de- 
composes into  antimony  and  hydrogen,  even  at  —  56°.  From  this 
instability  it  follows  that  stibine  cannot  be  obtained  free  from 
hydrogen  at  ordinary  temperatures.  Stibine  changes  to  a  liquid  at 
-91°.5;  the  liquid  boils  at  —18°,  and  solidifies  at  — 102°.5,  form- 
ing a  snow-like  mass ;  when  passed  through  a  heated  glass  tube  it 
decomposes  into  antimony  and  hydrogen,  just  as  arsine  does  into 
arsenic  and  hydrogen ;  the  antimony  is  deposited  in  the  form  of  a 
mirror,  which  is  more  metallic  in  appearance  than  is  the  one  formed 
of  arsenic,  and  which,  furthermore,  is  not  readily  volatilized.  When 
ignited,  stibine  burns  with  a  white  flame,  giving  off  a  dense  smoke 
of  the  trioxide  of  antimony  ;  a  spot  of  metallic  antimony  will  form 
on  a  cold  porcelain  plate  held  in  the  flame.  Of  course  chlorine, 
bromine,  or  iodine  attacks  stibine,  producing  the  corresponding 
halhydric  acid  and  the  halogen  compound  of  antimony. 

As  the  basic  properties  of  the  hydrogen  compounds  of  the  ele- 

*  The  alloy  consists  of  three  parts  zinc  to  two  of  antimony. 

v 

Of  THS 


252  ANTIMONY;  HALIDES  OF. 

merits  of  this  family  have  already  disappeared  when  arsine  is 
reached,  it  follows  that  stibine  can  form  110  salts  corresponding  to 
those  of  ammonium  and  phosphonium. 

Another  (solid)  compound  of  antimony  and  hydrogen  has  been 
described ;  to  this  the  formula  Sb4  H  has  been  ascribed ;  its  exist- 
ence, however,  is  doubtful. 

Antimony  forms  two  series  of  compounds  with  the  halogens, 
Sb  X3  and  Sb  X5 ;  they  correspond  to  those  of  phosphorus. 

SbF3,  solid. 

Sb  C13 ,  solid ;  soft  crystalline  mass,  melts  73°,  boils  223°. 

Sb  Br3 ,  solid ;  melts  90°,  boils  280°. 

Sb  I3 ,  solid ;  melts  167°,  boils  401°. 

Sb  F5  ,  not  crystalline,  gum-like  mass. 

Sb  C15 ,  fluid;  melts  — 6°,  can  be  distilled  without'decomposition  in  a  vacuum; 
boils  at  68°,  14  m.m.  pressure. 

Sb  I5 ,  solid ;  melts  78°,  and  decomposes  when  beated  to  a  higher  tempera- 
ture. 

The  trihalogen  compounds  can  be  produced  by  direct  union  of 
the  elements,  and  the  pentahalogen  compounds  by  addition  of  halo- 
gen to  the  trihalides.  The  trichloride  has  been  most  thoroughly 
studied,  and  its  chemical  behavior  will  serve  as  an  example  for  that 
of  the  other  haloid  compounds.  The  trichloride  is,  as  would  be  ex- 
pected, much  less  readily  decomposed  by  water  than  the  correspond- 
ing trichloride  of  arsenic ;  indeed,  it  can  even  be  dissolved  in 
tolerably  dilute  hydrochloric  acid  without  decomposing.  When 
ifc  is  added  to  water,  it  does  not  completely  break  down,  but  changes 
into  a  so-called  basic  salt. 

A  basic  salt  is  one  which  is  formed  by  the  interaction  of  less  of 
the  acid  and  more  of  the  base  than  is  necessary  to  produce  the  normal 
salt. 

We  can  consider  all  oxides  of  the  metals  which  act  as  bases  as 
being  derived  from  the  corresponding  hydroxides  by  loss  of  water, 
the  manner  of  formation  being  similar  to  that  of  the  anhydrides :  — 

H2S04-     H20  =  S03;        Ca(OH)2-     H20  =  CaO; 
2P  (OH)8  -  3  H20  =  P203;  2Fe  (OH)3  -  3  H20  =  Fe203. 

With  this  in  mind  we  can  define  a  basic  salt  as  one  in  which 
only  a  portion  of  the  hydroxyl  groups  of  a  base  have  reacted  with 
an  acid  to  form  a  salt  and  water ;  it  is  therefore  the  reverse  of  a 
so-called  acid  salt  (page  140  and  foot-note),  in  which  only  a  portion 


ANTIMONY  ;   BASIC    SALTS   OF.  253 

of  the  hydrogen  in  an  acid  has  been  replaced  by  a  metal.  The 
same  nomenclature  can  therefore  be  adopted  with  both  classes  of 
salts,  and  the  term  primary,  secondary,  and  tertiary  basic  salts  can 
be  used.  When  water  is  added  to  the  trichloride  of  antimony  the 
following  reaction  takes  place  :  — 

f  Cl  +  HOH  (  OH 

Sb-3  Cl  +  HOH  =      Sb-5  OH  +  2HC1. 

(  Cl  (  Cl 

The  compound  Sb  ( OH  )2  Cl  is  therefore  the  primary  basic 
chloride  of  antimony ;  the  two  hydroxyl  groups  contained  therein 
can  afterward  separate  water,  as  follows  :  — 

(OH 

Sb^OH= 
(Cl 

The  compound  Sb  0  Cl  is  therefore  also  a  basic  salt.  In  this 
case  the  group  of  elements  Sb  0  is  a  radicle  which  chemically  re- 
sembles a  monovalent  metal,  as  a  comparison  of  the  following 
formulae  will  make  clear  :  — 

(SbO)Cl,* 
Nad, 
KC1. 

The  monovalent  radicle  Sb  0  is  sometimes  called  stibionyl,  and 
is  quite  frequently  encountered  in  the  basic  salts  of  antimony ;  such 
an  instance  is  found  in  the  formula  of  stibionyl  sulphate  (Sb  0)2  S04. 

Those  metals  which  have  not  a  very  pronounced  metallic  char- 
acter are  the  ones  which  form  basic  salts;  metals  like  sodium, 
potassium,  or  calcium  do  not  so  produce  them. 

Antimony  trichloride  can  combine  with  the  chlorides  of  a  num- 
ber of  metals  to  form  double  chlorides  with  formulae  like  the  fol- 
lowing, Sb  C13 ,  3  K  Cl ;  more  extended  mention  of  these  will  be 
made  in  the  chapter  on  aluminium. 

The  pentachloride  of  antimony  is  an  unexpectedly  stable  com- 
pound ;  it  can  be  boiled  in  a  vacuum  without  change,!  and,  further- 
more, it  is  not  decomposed  by  cold  water,  but  forms  a  crystalline 
substance  containing  water  of  crystallization. 

*  More  complicated  basic  chlorides  than  this  one  also  exist;  for  their 
study  the  student  must  refer  to  a  larger  work. 

t  Anschiitz  and  Evans;  Liebig's  Annalen;  239,  285. 


254  ANTIMONY  ;    TRIOXIDE. 


CHAPTER   XXXIV. 

THE    COMPOUNDS    OF    ANTIMONY    WITH    OXYGEN,    AND    WITH 
OXYGEN    AND  HYDROGEN.     THE   SULPHIDES   OF   ANTIMONY. 

ANTIMONY  forms  the  following  oxides,  Sb203,  Sb204,  Sb205; 
only  the  last  one  of  these  is  a  pronounced  acidic  anhydride ;  indeed, 
the  trioxide  acts  as  a  base  when  brought  in  contact  with  pronounced 
acids. 

Antimony  trioxide  is  found  in  nature  as  the  mineral  senarmon- 
tite.  It  can  be  formed  by  burning  antimony,  or  by  oxidizing  the 
element  with  dilute  boiling  nitric  acid.  Chemically,  it  acts  just 
like  the  corresponding  oxide  of  arsenic ;  it  dissolves  in  alkaline 
hydroxides  to  form  meta-antimonites  :  — 

Sb203  +  2  KOH  =  2  Sb(XNa  +  H20. 

Two  hydroxides  derived  from  this  oxide  are  known ;  one,  which 
might  be  called  pyro-antimonous  acid,  has  the  formula,  Sb2  05  H4 ; 
the  other  is  the  normal  hydroxide,  Sb(OH)3,  corresponding  to 
orthophosphorous  acid,  P  (OH  )3 .  Antimony  trioxide  dissolves  in 
pronounced  acids  to  form  salts  in  which  antimony  is  the  metal ;  so, 
for  instance,  it  produces  the  trichloride  with  hydrochloric  acid  :  — 

Sb203  +  6H  Cl=  2  Sb  C13  +  3H20. 

The  salts  formed  with  other  than  the  halhydric  acids  are  basic 
ones,*  insoluble  in  water ;  while  the  trichloride,  formed  by  the 
above  reaction,  is  also  converted  into  the  insoluble  basic  chloride  by 
adding  water  to  the  acid  solution  (see  page  252).  Antimony  triox- 
ide is  easily  reduced  to  the  metal  by  heating  with  charcoal  or  in 
hydrogen.  When  heated  in  the  air  it  takes  up  oxygen  and  changes 
to  the  tetroxide,  Sb2  04 ,  which  substance  is  likewise  formed  from 
the  pentoxide  by  heating,  so  that  Sb2  04  is  the  most  stable  of  the 
oxides  of  antimony.  The  trioxide  is  but  little  soluble  in  water. 

TJie  pentoxide  of  antimony,  Sb205,  is  produced  by  oxidizing  anti- 
mony, with  fuming  nitric  acid  (page  202),  or  with  aqua  regia  (page 

*  A  neutral  and  an  acid  sulphate  have  been  described  by  Adie ;  Journ. 
Chem.  Soc. ;  1890  ;  540. 


ANTIMONIC    ACIDS.  255 

203)  ;  it  loses  oxygen  at  red  heat  and  changes  into  the  tetroxide, 
Sb204. 

The  basic  properties  of  Sb203  have  been  entirely  destroyed  by 
adding  the  two  oxygen  atoms  necessary  to  form  Sb2  05 ,  so  that 
the  latter  substance  forms  no  salts  with  acids. 

The  following  acids  derived  from  antimony  pentoxide  are  known ; 
they  correspond  to  those  of  phosphorus  and  of  arsenic,  and  are :  — 
Sb  03H,  meta-antimonic  acid, 

Sb  04H3,  ortho-antimonic  acid, 

Sb207H4,  pyro-antimonic  acid, 

An  acid   having  the  formula,  Sb2  09  H8,  is  also  known ;    this  is  a 
hydrated  pyro-antimonic  acid  formed  as  follows :  — 

Sb207H4  +  2H20  =  Sba09H8. 

The  acid  H8  Sb2  09  changes  into  ortho-antimonic  acid  on  standing 
with  water ;  it  is  converted  into  pyro-antimonic  acid  at  100° ;  pyro- 
antimonic  acid  forms  meta-antimonic  acid  at  200° ;  and  this  is  finally 
converted  into  the  anhydride  Sb205  at  300°. 

Only  two  sulphides  of  antimony,  Sb2S3  and  Sb2S5,  are  known 
with  certainty.  They  are  entirely  analogous  to  the  corresponding 
arsenic  compounds. 

The  trisulphide,  Sb2  S3 ,  has  already  been  mentioned  as  the  min- 
eral antimonite.  On  roasting  in  a  current  of  air  it  is  changed, 
first  into  the  trioxide,  Sb203,  and  then  into  the  tetroxide,  Sb204. 
The  amorphous  sulphide  is  produced  by  the  addition  of  hydrogen 
sulphide  to  an  acidified  solution  of  antimony  trioxide,  just  as  the 
corresponding  arsenic  compound  is  formed.  It  is  dark  orange- 
colored  powder,  insoluble  in  water  or  dilute  acids,*  soluble  in 
alkaline  hydroxides  or  sulphides ;  by  disolving  in  the  latter  it  forms 
a  pyro-sulphantimonite ;  in  this  way  antimony  differs  from  arsenic, 
which  produces  the  meta-sulpharsenite  under  similar  circumstances :  — 

Sb2  S3  +  4  KSH  =  Sb,  S5  K4  -f  2  H2  S.f 

(The  reaction  for  the  formation  of  antimonite  and  sulphantimonite 
by  dissolving  Sb2  S3  in  alkaline  hydroxides  is  given  under  arsenic, 

*  Under  certain  conditions  antimony  trisulphide  can  be  obtained  in  a  form 
which  is  soluble  in  water.  Schulze  ;  Journal  fiir  Praktische  Chemie  ;  [2]  27, 
320.  The  statement  that  a  meta-sulphantimonite  is  also  produced  is  made  Ly 
Ditte;  Compt.  Rend.;  102,  168. 

t  Compare  with  the  following  :  — 

Sb2  O3  +  6  KOH  =  Sb  03  K3  +  3  H2  O. 


256  ANTIMONY  ;    PENTASULPHIDE. 

page  245,  foot-note.)  When  the  alkaline  sulphide  contains  dissolved 
sulphur,  the  sulphaiitimonate  is  produced ;  the  same  *  is  the  case  with 
arsenic. 

Antimony  pentasulphide  can  be  precipitated  by  adding  sulphu- 
retted hydrogen  to  a  solution  containing  antimonic  acid ;  f  it  bears 
a  complete  resemblance  to  the  corresponding  arsenic  compound,  dis- 
solving in  alkaline  sulphides  or  hydroxides  to  form  sulphantimon- 
ates,  the  majority  of  which  are  salts  of  the  ortho  acid,  H3SbS4. 
Only  the  salts  of  the  alkalies  and  alkaline  earths  $  are  soluble  in 
water.  The  pentasulphide  of  antimony  is  precipitated  and  hydro- 
gen sulphide  is  formed  when  acids  are  added  to  solutions  contain- 
ing the  sulphantimonates. 

It  is  scarcely  necessary  to  add  a  table  of  the  antimony  compounds, 
as  they  correspond  so  completely  to  those  of  arsenic ;  one  fact  can  be 
reiterated,  namely,  that  antimony  can  form  pentahalogen  derivatives, 
while  arsenic  is  unable  to  do  so. 

*  Cl.  Zimmerman  (personal  information). 

t  Difference  between  antimony  and  arsenic.    (See  page  246  and  foot-note. ) 

J  The  alkaline  earths  are  calcium,  barium,  and  strontium. 


BISMUTH  ;   OCCURRENCE,  PREPARATION.  257 


CHAPTEE   XXXV. 

BISMUTH.  THE  COMPOUNDS  OF  BISMUTH  WITH  THE  HALO- 
GENS, WITH  OXYGEN,  WITH  OXYGEN  AND  HYDROGEN,  AND 
WITH  SULPHUR. 

Bismuth ;  Symbol,  Bi ;  Atomic  iveight,  208.9 ;  Specific  gravity 
(electrolytic),  9.74;  specific  gravity  of  vapor,  air=l,  is  10.1,  H2 
=2,  is  290.8 ;  molecular  weight  of  Bi2  is  417.8. 

BISMUTH,  the  element  having  the  highest  atomic  weight  in  this 
family,  is,  as  would  be  expected,  entirely  a  metal,  although  its  com- 
pounds correspond  in  formula  to  those  of  the  other  members. 

The  element  occurs  in  nature  as  :  — 

Native  bismuth,  accompanying  cobalt,  nickel,  and  lead  ores  in  gneiss  and 
other  crystalline  rocks  and  in  clay  slate. 

Bismuth  trioxide,  Bi2O3;  an  earthy  mass  called  bismite. 
Bismuth  trisulphide,  Bi2  S3 ;  bismuthinite. 
Bismuth  telluride,  Bi2  (  Te,  S)8  ;  tetradymite. 

It  has  not  been  ascertained  with  certainty  how  early  in  the  his- 
tory of  metallurgy  bismuth  became  known ;  it  was  formerly  con- 
fused with  a  number  of  minerals,  all  of  which  went  by  the  name  of 
marcasite.  It  first  became  known  as  a  new  metal  in  the  fifteenth 
century,  and  was  called  bisemat  or  ivisemutum,  but  was  even  then 
confounded  with  antimony ;  at  the  end  of  the  eighteenth  century, 
however,  it  was  universally  considered  to  be  a  metal. 

Bismuth,  like  antimony,  is  prepared  from  its  ores  by  roasting 
the  sulphide  with  iron  or  the  oxide  with  charcoal.  The  metal  which 
finds  its  way  into  commerce  contains  small  quantities  of  arsenic  and 
iron;  the  natural  metal  is  nearly  chemically  pure.  The  bismuth 
which  is  to  be  used  in  the  preparation  of  compounds  intended  for 
pharmaceutical  purposes  is  subjected  to  special  processes  of  purifi- 
cation, for  all  arsenic  must  be  removed  from  it. 

Bismuth  is  a  reddish  metal  with  a  pronounced  metallic  lustre 
and  a  coarse  crystalline  structure.  Its  specific  gravity  is  9.74.  The 
metal  is  brittle,  neither  malleable  nor  ductile,  and  can  readily  be 


258  BISMUTH  ;    PROPERTIES  ;  HALIDES   OF. 

pounded  into  a  powder.  It  melts  at  270°,  and,  in  solidifying,  ex- 
pands just  as  water  does  in  forming  ice.  It  boils  at  about  1400° ; 
the  specific  gravity  of  its  vapor  at  white  heat,  (1600° -1700°)  is 
10.1.*  The  calculated  specific  gravity,  were  the  molecule  Bi2,  is 
14.4,  and  is  7.2  if  the  same  were  composed  simply  of  individual 
atoms ;  as  the  number  10.1  is  considerably  less  than  that  which 
would  be  found  were  the  molecule  diatomic  and  to  be  expressed  by 
the  formula  Bi2,  it  follows  that  the  vapor  of  bismuth  consists,  in 
part  at  least,  of  the  individual  atoms.  At  a  slightly  higher  temper- 
ature than  that  of  the  experiment,  the  molecular  and  atomic  weights 
of  bismuth  would  probably  be  identical. 

Dilute  hydrochloric  or  sulphuric  acids  do  not  attack  bismuth ; 
concentrated  hydrochloric  acid  has  but  little  effect ;  hot  and  concen- 
trated sulphuric  acid  dissolves  it,  forming  the  sulphate  of  bismuth 
and  sulphur  dioxide  (see  page  137)  ;  nitric  acid  quite  readily  forms 
the  nitrate,  while  it  is  itself  reduced.  (See  page  206.) 

A  number  of  alloys  of  bismuth  are  commercially  important ; 
they  all  have  low  melting  points ;  for  example,  a  mixture  of  five 
parts  of  bismuth,  three  of  lead,  two  of  tin,  and  three  of  cadmium 
fuses  at  65°.  One  of  the  chief  uses  of  these  alloys  is  in  the  copy- 
ing of  wood-cuts  and  in  stereotyping. 

The  halogens  form  the  following  compounds  with  bismuth :  — 

Bi  F3 ,  bismuth  trifluoride. 

Bi  C13 ,  "bismuth  trichloride. 

Bi  Br3 ,  bismuth  tribromide. 

Bi  I3 ,  bismuth  tri-iodide. 

All  of  these  compounds  are  solid  bodies  at  ordinary  temperatures.  Other 
halogen  derivatives  of  bismuth  are  described ;  they  are  unstable  bodies  which 
have  probably  not  been  obtained  in  a  pure  state. 

The  trihalogen  compounds  can  either  be  formed  by  the  direct 
union  of  bismuth  with  chlorine,  bromine,  or  iodine,  or  by  dissolving 
the  trioxide,  Bi203,  in  concentrated  halhydric  acids,  this  oxide  of 
bismuth  being  entirely  basic  in  its  character.  The  chloride  or 
bromide  of  bismuth  can  be  dissolved  in  very  little  water  or  in  acids 
without  change ;  these  salts,  however,  decompose  on  the  addition  of 
an  excess  of  water  just  as  do  the  corresponding  antimony  com- 
pounds (see  page  253),  while  at  the  same  time  the  insoluble  basic 
chloride  or  bromide  is  formed  :  — 

*  H.  Biltz  and  V.  Meyer;  Ber.  d.  Deutsch.  Chem.  Gesell.;  22,  725. 


BISMUTH;    PROPERTIES;    OXIDES   OF.  259 

(  Cl  (OH 

BiJ  C1  +  2H20  =      BiJ  OH  +  2HC1, 
(01  (01 

(OH 

BiJOH=      B^Xi  -fH20. 
(01 

The  iodide  of  bismuth  is  decomposed  only  by  boiling  water ;  it  is 
insoluble  in  cold  water.  The  halogen  compounds  of  bismuth  form 
a  number  of  double  salts  which  will  be  discussed  later  j  *  the  for- 
mulae of  a  few  of  these  salts  are  as  follows :  — 

K  Cl,  Bi  C13  +  H2  O  ;  2  Na  Cl,  Bi  C13  +  H2  O. 

2  K  Cl,  Bi  C13  +  2  H2  O  ;  2  Na  Cl,  Bi  C13  +  3  H2  O. 

An  acid  composed  of  the  fluoride  of  bismuth  and  of  hydrofluoric  acid, 
Bi  F3 ,  3  HF  is  also  known. 

The  following  oxides  of  bismuth  are  known :  — 

Bi  O,  bismuth  monoxide,  brownish  black,  crystalline  powder, 
Bi2O3,  bismuth  trioxide,  light  yellow  powder, 
Bi2  O4 ,  bismuth  tetroxide,  dark  brown  powder, 
Bi2O5,  bismuth  pentoxide,  brown  powder. 

None  of  these  oxides  dissolve  in  alkaline  hydroxides  to  form 
salts.  The  one  with  least  oxygen,  Bi  0,  is  produced  by  the  reduc- 
tion of  salts  derived  from  the  oxide  Bi2  O3 ;  f  it  is  insoluble  in 
water,  and,  when  dry,  oxidizes  so  readily  in  the  air  that  the  com- 
pound burns  like  a  piece  of  tinder ;  the  trioxide  is  produced  by  this 
combustion. 

Bismuth  burns  at  a  high  red  heat,  forming  the  trioxide,  Bi2  03 . 
This  substance  can  be  more  readily  prepared  by  the  addition  of 
potassium  hydroxide  to  a  solution  of  a  bismuth  salt.  The  first 
product  formed  is  the  hydroxide,  Bi  (OH)3,  which  is  insoluble  in 
water :  -  Bi  (No3)8  +  3  KOH  =  Bi  (OH)3  +  3  KNO,. 

This  reaction  is  a  usual  one  with  most  metals,  for  by  far  the 
greater  number  of  metallic  hydroxides  are  insoluble  in  water.  The 
following  reactions  are  typical :  — 

*  For  a  discussion  of  the  nature  of  double  salts,  see  chapter  on  Aluminium, 
and  also  Remsen;  Amer.  Chem.  Journ.;  14,  165,  and  Inorganic  Chemistry 
(Holt  &  Co.,  1890),  page  461. 

t  By  treating  solution  of  chloride  of  bismuth,  in  hydrochloric  acid,  with 
an  alkaline  solution  of  stannous  chloride,  Sn  C12 .  See  chapter  on  Tin. 


260  BISMUTH   NITRATE,    SUBNITRATE. 

Fe  S04  +  2  KOH  =  K2  S04  +  Fe  (OH)2, 
Cu  S04  +  2  KOH  =  K2  S04  +  Cu  (OH),, 
Mg  (N03)2  +  2  KOH  =  K2  (N08)2  +  Mg  (OH)2 . 

In  each  of  these  cases  the  insoluble  hydroxide  is  precipitated  by 
the  addition  of  a  soluble  hydroxide  to  the  salts  of  the  metals.  The 
hydroxides  of  the  alkalies  are  the  only  ones  which  are  readily  solu- 
ble in  water ;  those  of  the  metals  calcium,  barium,  and  strontium  are 
soluble  with  some  difficulty,  all  others  are  insoluble.  Those  hy- 
droxides of  the  metals  which  are  not  capable  of  being  dissolved  by 
water  are  generally  converted  into  the  corresponding  oxides  by  the 
application  of  a  very  moderate  heat.  The  hydroxide  of  bismuth 
loses  water  below  100°,  and  is  converted  into  a  compound  having 
the  formula  Bi  02  H :  *  — 

(OH 

Bi^OH  =     Bi 
(OH 

and  this,  lastly,  above  boiling  heat,  is  changed  to  the  trioxide :  — 

2  Bi  02  H  =  Bi,  03  +  H2  0, 
just  as  nitrous  acid  changes  into  water  and  N203. 

The  various  salts  of  bismuth  are  produced  by  dissolving  the 
oxide  or  the  hydroxides  in  acids.  The  chloride  has  already  been 
discussed,  the  only  other  ones  we  need  mention  are  the  nitrate  and 
sulphate.  . 

BISMUTH  NITRATE.  —  Bi  (  NO3  )8  +  5£  H2  O.t  Prepared  by  dissolving 
either  the  element  bismuth,  or  the  oxide,  or  one  of  the  hydroxides  in 
nitric  acid.  Bi2  O3  +  6  HNO3  =  2  Bi  (  NO3  )3  +  3  H2  O.  Clear,  large 
crystals,  which  melt  at  80°  in  their  water  of  crystallization,  and  which 
then  give  off  nitric  acid,  forming  the  basic  nitrate  at  120°. 

BASIC  BISMUTH  NITRATE.  —  (  Subnitrate  of  bismuth  )  Bi  ONO3  +  H2  O,f 

or  (OH 

Bi^OH 

(N03. 

Formed  by  decomposing  the  nitrate  with  hot  water.  Insoluble  in 
water.  It  is  a  primary  basic  salt  in  which  one  hydroxyl  group  has 
been  replaced  by  the  group  NO3;  when  heated  in  a  platinum  dish,  it 

*  This  hydroxide  is  the  one  corresponding  to  nitrous  acid  NO2  H. 
t  Yvon;  Bull.  Soc.  Chim. ;  [2],  27,  491. 

J  More  complicated  formulae  have  been  assigned ;  for  these  a  larger  man- 
ual must  be  consulted. 


BISMUTH   PENTOXIDE  ;    SULPHIDES.  261 

changes  to  the  oxide  Bi2  O3 ,  giving  off  nitric  acid.  The  salt  Bi  ONO3 
can  be  considered  as  derived  from  the  hydroxide  Bi  O2  H,  by  replacing 
the  hydroxyl  group  with  NO8 . 

Bi  |  °H  +  HN03  .  Bi  <[  °n_  +  H2  O. 


The  subnitrate  of  bismuth  is  the  most  important  salt  of  bismuth, 

being  extensively  used  in  medicine. 

BISMUTH  SULPHATE.  —  Bi2  (SO4  )3  +  7  H2  O.  Prepared  by  dissolving 
the  oxide  in  concentrated  sulphuric  acid  and  then  adding  water.  At 
100°  it  changes  to  Bi2  (SO4  )3  +  3  H2  O.  Boiling  water  changes  it  into 
the  insoluble  basic  sulphate. 

fOBiO 

0  J  OBiO  =  (BiO)2S04. 

01  o 

to 

The  univalent  group  Bi  O  corresponds  to  stibionyl  (page  253). 

The  pentoxide  of  bismuth,  Bi2  05 ,  is  produced  by  oxidizing  the 
trioxide  by  means  of  hydrogen  peroxide.  It  is  an  orange  or  brown 
colored  powder,  which  forms  but  one  hydroxide,  the  latter  corre- 
sponding to  the  meta-acids.  This  hydroxide,  Bi  03  H,  is  changed  to 
the  oxide  Bi2  05  on  heating.  Neither  the  oxide  nor  hydroxide  has 
basic  or  acid  properties.* 

Bi2  04  has  been  considered  as  a  compound  formed  by  replacing 
the  hydrogen  of  Bi  03  H  by  the  univalent  group  of  elements  Bi  0, 
acting  as  a  metal :  — 

Bi  0  Bi  03  =  Bi2  04 . 

The  sulphides  of  bismuth  correspond  to  the  oxides  Bi  0  and 
Bi2  03 .  Bismuth  monosulphide,  Bi  S,  is  of  little  importance.  Bis- 
muth trisulphide  is  found  in  nature  as  the  mineral  bismuthinite. 
It  can  be  produced  by  fusing  bismuth  and  sulphur  together,  or  it 
can  be  formed  by  the  action  of  hydrogen  sulphide  on  a  solution  of 
a  bismuth  salt :  — 

2  Bi  C13  +  3  H2  S  =  Bi2  S3  +  6  H  Cl. 

The  sulphide  of  bismuth,  like  the  oxide,  has  no  resemblance  in 
its  action  to  the  acidic  anhydrides ;  it  cannot  be  dissolved  in  the 
hydroxides  or  sulphides  of  the  alkalies  to  form  sulpho-salts.  On 
the  other  hand,  it  is  readily  attacked  by  concentrated  hydrochloric 
acid  to  form  the  trichloride  of  bismuth  and  hydrogen  sulphide  :  — 


262 


ELEMENTS    OF   NITROGEN   FAMILY ; 


Nitric  acid  dissolves  it  to  form  the  nitrate  of  bismuth ;  while  the 
hydrogen  sulphide  which  is  set  at  liberty  is,  of  course,  oxidized  by 
the  excess  of  nitric  acid  to  form  sulphur,  and,  finally,  sulphuric 
acid. 

The  elements  of  the  nitrogen  family,  as  we  have  seen,  show  the 
same  graduation  in  properties  with  increasing  atomic  weights  as  was 
displayed  by  the  elements  belonging  to  the  sulphur  group.  This  is 
best  brought  to  light  by  the  following  table  :  — 


ATOMIC 
WEIGHTS. 

USUAL  PHYSICAL 
CONDITION. 

MELTING 
POINTS. 

BOILING 
POINTS. 

SPECIFIC 
GRAVITY. 

PBOPEETIES. 

N      14.03 

A  gas. 

—  203° 

—  193° 



not-metallic. 

P      31. 

Yellow,  solid,  eas- 
ily fused.* 

44° 

250° 

1.82 

As    75. 

Steel  gray,  brittle, 
crystalline. 

About  450°  t 

800"  t 

5.7 

Sb  120. 

Silver  white,  brit- 
tle, crystalline. 

About  425" 

About  1300° 

6.7 

Bi  208.9 

Reddish,  metallic, 
brittle,  crystal- 
line. 

270° 

About  1600° 

9.8 

metallic. 

SPECIFIC   GRAVITY   OF   VAPORS. 


AIR  =  1. 

H2  =  2. 

TEMP. 

MOLECULES. 

N. 

.9713 

27.9 

Ordinary 

X2 

P. 

4.16 

119.8 

800° 

P4 

3.14 

90.4 

1708° 

P2  and  P4 

As. 

10.3 

296.6 

900° 

As4 

5.37 

154.0 

1736° 

As2 

Sb. 

9.78 

281.6 

1640° 

Sb2 

Bi. 

10.1 

290.8 

1650° 

Bi!  and  Bi2 

*  Red  phosphorus  in  part  is  then  changed  to  yellow  at  261°.  Specific 
gravity,  2.08  to  2.14. 

t  Arsenic  volatilizes  without  previously  melting,  unless  it  is  under  pres- 
sure. 


COMPARATIVE    TABLE    OF    COMPOUNDS.  263 

HYDROGEN   COMPOUNDS  XH3 ,   X2  H4 ,    X3  H. 


HEAT 

OF 

STABIL- 

BASIC PBOPERTIES. 

FORMA- 

ITY. 

TION. 

N 

NH3 

120  K 

NaH4 

N3H 

NH3  +  HI  =  NH4  I  \ 

The   stability 

P 

PH3 

43  K 

Pa  H4 

P4H2 

\ 
PH3+HI=PH4I    \ 

and  basic  prop- 
e  r  t  i  e  s       d  i- 

\ 

m  i  n  i  s  h    with 

As 

As  H3 

-441  K 

As~H, 

As  H3  

increasing    me- 

Sb 

^b  H 

Sb  H 

tallicproperties 
of    elements; 

only    ammonia 

Bi 

combine     with 

acids    to    form 

salts.* 

CHLORINE  COMPOUNDS,  X  C13 ,  XC15. 


N 

N    C13 

Liquid,  explosive, 

t  Completely    decomposed 

P 

P    C13 

"      decomposed  by  H2  O  t 

PC15 

Solid 

by  water  as  follows  :  — 
X  C13  +  3  H2  O  =  X  (O  H  )3 

As 

As  Cl 

«                                i«                           «                 «          4. 

•+•  3  H  Cl. 

Sb 

^\.S   V^13 

SbCl3 

Solid,            "            »        "    t 

SbCl6 

Solid. 

$  Partially  decomposed  by 
water,  forming  basic   chlo- 

Bi 

Bi  Cl 

U                                    <(                            («                 (C          4- 

»•  '   1      •   o      -T    11 

1>1    V^13 

riUtrS  US  lOUOWS  .  

XC13  +  2H2O  =  X(OH)2 

Cl  ;  X  (OH  )2  Cl  =  X  O  Cl  + 

H2O. 

The    pentachlorides    with 

water   yield    phosphoric   or 

antimonic  acid. 

OXIDES,  HYDROXIDES,  SULPHIDES,  AND  SULPHO  SALTS. 


OXIDES.       X2O,  X2O3,X2O6. 

SULPHIDES. 

N2O 

N    O 
N203 
N204 
N206 

BiO 

Bi203 
Bi204 
Bi205 

As2  S2 
As2S3 

As2O3 

1%03 
P»04 

P206 

Sb203 
Sb204 
Sb206 

F2S3 

Sb2S3 

Bi2S3 

As205 

1%SB 

As2S6 

Sb2SB 



*  Phosphine  forms  only  a  limited  number  of  phosphonium  salts;  these 
are  decomposed  by  water.  The  compound  N2  H4  forms  salts  with  acids ;  P2  H4 
does  not.  N3  H  is  acid  in  its  character,  resembling  H  Cl,  H  Br,  and  H  I. 


264  ELEMENTS    OF   NITROGEN    FAMILY. 

The  acids  derived  from  the  oxides  X.2  O3  are  formed  according  to  the  types 
X  (OH)3,  ortho  acids;  XO2H,  meta  acids;  H4X2O5,  pyro  acids.  Nitrogen 
forms  only  the  meta  acids  and  salts  derived  from  this ;  phosphorus  forms  only 
the  ortho  acids,  but  salts  of  the  two  other  ones  are  known ;  arsenic  forms 
no  hydroxides,  the  salts  of  the  meta  arsenious  acid  are  the  most  frequent; 
antimony  forms  the  hydroxides  corresponding  to  the  ortho  and  pyro  acid, 
the  meta-antimonites  are  the  most  frequent;  bismuth  forms  the  hydroxide 
Bi(OH)3,  it  has  no  acid  properties.  The  trioxides  of  nitrogen  and  phos- 
phorus are  acidic  only,  those  of  arsenic  and  antimony  are  both  basic  and 
acidic,  that  of  bismuth  is  basic  only. 

The  pentoxides  are  all  acidic  with  the  exception  of  that  of  bismuth,  which 
is  neither  acidic  nor  basic ;  the  acids  derived  from  the  oxides  X2  O5  are  formed 
according  to  the  types  XO4H3,  ortho  acids  ;  XO3H,  meta  acids;  X2O7H4, 
pyro  acids.  Nitrogen  forms  only  the  meta  acid,  NO3H;  bismuth  forms  one 
hydroxide  Bi  O3  H,  not  acid  in  its  nature. 

The  sulphides  As2S8,  As2S5,  Sb2S3,  Sb2  S5  are  insoluble  in  water;  they 
dissolve  in  alkaline  sulphides  to  form  salts  of  sulpho  acids.  These  acids  are 
exactly  like  the  oxygen  acids  with  the  exception,  that  in  them  sulphur  has 
taken  the  place  of  oxygen,  atom  for  atom.  The  sulphides  P2  S3 ,  P2  S5 ,  are 
decomposed  by  water;  the  sulphide  Bi2  S3  has  no  acidic  properties. 


ELEMENTS    OF   CARBON   FAMILY.  265 


CHAPTER   XXXVI. 

THE   ELEMENTS   OF   THE   CARBON   FAMILY. 

THE  elements  of  the  carbon  family  are  carbon,  silicon,  germa- 
nium, tin,  and  lead.  Changes  similar  to  those  observed  in  the  pre- 
ceding family  are  caused  by  the  increase  in  the  atomic  weights 
belonging  to  the  elements  in  this  one ;  but,  as  the  whole  family  is 
less  not-metallic  than  is  that  of  which  nitrogen  is  the  representa- 
tive, only  two  of  the  elements,  namely,  carbon  and  silicon,  can  form 
hydrogen  compounds ;  the  transition  from  not-metal  to  metal  takes 
place  after  the  second  member  of  the  group. 

The  alterations  in  the  physical  properties  of  the  isolated  ele- 
ments show  this  increasing  metallic  character  ;  for  while  carbon 
and  silicon  are  found  either  in  the  forms  of  amorphous  black 
elements  or  crystalline  bodies  which,  in  the  case  of  diamond,  may 
even  be  transparent,  germanium,  tin,  and  lead  have  a  brilliant 
metallic  lustre.  Germanium  and  tin,  however,  have  a  crystalline 
structure,*  while  lead  is  the  only  member  of  the  family  which  is 
perfectly  malleable  and  ductile.  The  fusing  points  of  the  elements 
under  discussion  diminish  with  increasing  atomic  weights,  while  the 
specific  gravities  increase. 

Carbon;  infusible,  possibly  softens  in  the  heat  of  the  electric  arc;  specific 
gravity  (as  graphite)  2.2. 

Silicon;  fuses  in  the  heat  of  the  electric  arc;  specific  gravity  (graphitoi- 
dal)  2.49. 

Germanium;  melts  at  about  900°;  specific  gravity  5.46. 

Tin;  melts  at  230°;  specific  gravity  7.29. 

Lead;  melts  at  325°;  specific  gravity  11.44. 

The  general  formulae  for  the  hydrogen  compounds  of  the  ele- 
ments of  the  chlorine,  oxygen,  and  nitrogen  families  are  respect- 
ively, XH,  XHa,  and  XH3,  in  the  carbon  family  the  corresponding 
compounds  are  XH4 .  The  power  of  fixing  hydrogen  atoms  which  is 

*  Compare  with  arsenic,  antimony,  and  bismuth  in  the  preceding,  less 
metallic,  family. 


266 


ELEMENTS    OF   CARBON   FAMILY;     COMPARISON. 


possessed  by  any  one  atom  of  a  not-metal  is  exhausted  when  the 
number  four  is  reached ;  indeed,  no  elements  other  than  those  which 
we  have  considered  are  capable  of  forming  gaseous  hydrogen  com- 
pounds with  definite  formulae.*  It  follows  that  the  maximum  val- 
ence which  any  element  displays  toward  hydrogen  alone  is  four  (see 
pages  107,  108).  The  relationship  between  the  groups  of  elements 
and  the  formulae  of  the  hydrogen  compounds  becomes  more  apparent 
if  we  arrange,  in  the  order'  of  their  atomic  weights,  the  symbols  of 
those  individuals  which  we  have  studied,  reversing  the  order  ob- 
served in  the  table  on  page  175  :  — 


ELEMENTS. 

HYDROGEN  COMPOUNDS. 

RELATION  TO  ATOMIC  WEIGHTS. 

C       N       O        F 

CH4     NH3       OH2       FH 

As  we  pass  from  left  to  right, 

12       14       16       19 

the  atomic  weights  of  the  ele- 

Si       P        S       Cl 

SiH4    PH3       SH2       C1H 

ments   on    any  horizontal    line 
increase,  while  their  valence  to- 

28     31       32   35.5 

ward  hydrogen  diminishes.    No 

As      Se      Br 

AsH3    SeH2    BrH 

elements  with    atomic    weights 

75      79      80 

lying  between  those  of  any  two 

—      Sb      Te        I 

Sb  H3    Te  H2    I  H 

on     any    horizontal     line     are 
known.       These    elements     are 

120    125     127 

, 

therefore  a  section  of  that  table 

which    would    be    obtained   by 

arranging  all  of  the  elements  in 

the    order    of    their    increasing 

atomic  weights.    (See  page  17.) 

The  elements  given  on  the  above  table  are  the  only  ones  which 
are  capable  of  forming  gaseous  hydrogen  compounds.  The  hydro- 
gen compound  of  silicon  is  more  easily  decomposed  than  is  that  of 
carbon ;  for  the  rule  is  without  exception  that,  with  increasing  atomic 
weight,  in  any  given  family,  there  is  a  diminution  in  the  stability  of 
the  hydrogen  compounds  with  the  members  of  that  family. 

Carbon  possesses  in  the  most  eminent  degree  that  property  which 
we  observed  in  a  rudimentary  form  in  the  hydrogen  compounds  of 
nitrogen ;  namely,  the  element  can  form  an  almost  unlimited  number 
of  complicated  hydrogen  compounds  derived  from  a  nucleus  of 
carbon  atoms,  united  one  with  the  other,  just  as  the  two  nitrogen 
atoms  are  joined  in  hydrazin  (page  192).  As  many  as  sixty 
carbon  atoms  are  known  to  be  thus  united  in  a  long  and  simple 
chain,  while  the  variety  of  compounds  may  be  almost  indefinitely 

*  See  chapter  on  boron  and  its  compounds. 


ELEMENTS    OF    CARBON    FAMILY;    COMPARISON.  267 

increased  by  branching  side  chains,  or  by  the  formation  of  rings  of 
atoms,  each  of  which  can  serve  as  a  nucleus  for  further  substitution 
or  addition.  The  study  of  these  compounds  forms,  at  present,  a 
separate  branch  of  chemistry,  which  is  generally  termed  organic 
chemistry,  although,  of  course,  there  is  no  real  distinction  between 
this  and  so-called  inorganic  chemistry.  A  few  of  the  simpler  car- 
bon and  hydrogen  compounds  will  be  taken  up  in  the  course  of  this 
work. 

All  of  the  elements  of  the  carbon  family  form  dioxides,  X02 , 
and  all  but  silicon  are  capable  of  producing  a  monoxide,  XO.  The 
dioxides  of  carbon,  silicon,  germanium,  and  tin  have  the  character 
of  acidic  anhydrides ;  the  dioxide  of  tin,  however,  like  the  trioxide 
of  arsenic,  can  be  both  acidic  and  basic,  for  it  dissolves  as  well 
in  acids  as  in  bases  to  form  salts.  The  monoxide  of  carbon  is 
neither  acidic  nor  basic,  that  of  germanium  is  slightly  basic,  the 
other  monoxides  are  all  basic  in  their  character  and  form  well- 
defined  salts  with  acids. 

The  acids  derived  from  the  oxides  have  the  general  formulae 
H2  XO3  and  H4  X04 ;  those  of  the  first  class  being  the  meta-acids, 
those  of  the  second,  the  ortho-acids;  these  compounds  all  are 
unstable,  readily  losing  water  and  leaving  the  corresponding  anhy- 
dride X0.j ;  indeed,  it  is  doubtful  if  carbonic  acid  exists  at  all,  even 
in  aqueous  solution.  The  silicic  acids,  both  ortho  and  meta,  are 
changed  to  silicon  dioxide  when  heated ;  they  lose  a  large  amount 
of  water  even  when  dried  at  ordinary  temperatures,  so  that  the 
existence  of  hydrated  silicic  acids  of  definite  formula  is  doubtful. 
Germanium  dioxide  apparently  forms  no  hydrates ;  the  stannic 
acids,  both  ortho  and  meta,  are  completely  dehydrated  when  heated 
to  redness.  Lead  dioxide  has  no  acidic  properties. 

All  inorganic  carbonates  are  derived  from  a  meta-carbonic  acid, 
ILCOg;  only  a  few  organic  derivatives  of  ortho-carbonic  acid, 
H4  CO4  are  known.  Both  ortho-  and  meta-silicates  exist  and  form 
two  classes  of  frequently  occurring  minerals,  while  salts  of  much 
more  complicated  silicic  acids  (formed  by  the  separation  of  water 
between  two  or  more  formula  weights  of  the  ordinary  acids)  are 
quite  common.  Both  ortho-  and  meta-stannic  acid,  H4  Sn  04  and 
H2  Sn  03 ,  are  known ;  although  all  salts  are  derived  from  the 
latter  compound.  The  relationships  between  the  compounds  dis- 
cussed above  are  made  more  apparent  by  the  following  table :  — 


268 


ELEMENTS   OF   CARBON    FAMILY;    COMPARISON. 


OXIDES. 

META-ACID8. 

OKTHO-ACIUS. 

-M  ETA-SALTS. 

ORTHO-SALTS. 

C/-v 

HP    A     # 

4. 

Mr*  c\ 

O2 

Si  O2 
f^  _  /-\ 

H2Si03* 

H4Si  04* 

•2  *>    ^3 

M2  Si  03 

MPo  O    + 

M4  Si  O4 

SnO2 

H2SnO3 

H4SnO4 

2  <^e  U3  j 
M2  Sn  O8 

M4  Sn  O4 

"b  O2 

The  monoxides,  with  the  exception  of  that  of  carbon,  are  bases ; 
they  dissolve  in  acids  to  form  a  number  of  well-defined  salts,  which, 
when  compared  with  the  salts  formed  from  the  basic  lower  oxides  of 
the  elements  of  the  preceding  family,  show  a  similar,  though  not 
quite  so  well  marked,  tendency  to  change  into  basic  salts  on  the 
addition  of  water. 

The  following  sulphides,  corresponding  to  the  oxides,  have  been 
studied :  — 


Carbon  monosulphide,  CS  (?) 


Carbon  disulphide,  C  S2 . 
Silicon  disulphide,  Si  S2 . 


Germanium  monosulphide,  GeS;  Germanium  disulphide,  GeS2. 
Stannic  monosulphide,  SnS;  Stannic  disulphide,  Sn  S2  . 

Plumbic  monosulphide,  Pb  S;        — 

The  disulphide  of  silicon  is  too  unstable  to  enter  into  other  com- 
pounds ;•  indeed,  it  is  decomposed  even  by  the  moisture  of  the  air  ; 
but  the  other  disulphides  dissolve  in  the  sulphides  of  the  alkali 
metals  to  form  sulpho-salts  which,  in  formula,  correspond  to  the 
oxy-salts. 

The  chlorine  compounds  and  chlorine  and  oxygen  compounds 
are  formed  after  the  general  formulae  X  C12,  X  C14,  and  X  0  C12 . 
Of  course,  a  representative  of  each  of  these  classes  is  not  known  for 
every  element  in  the  family ;  the  most  important  ones  will  be  indi- 
vidually discussed  in  the  succeeding  chapters. 

*  Existence  as  acids  doubtful. 

t  Only  orthocarbonates  of  organic  compounds  are  known, 
t  More  thorough  investigation  of  germanium  salts  is  necessary. 
§  Carbon  also  forms  a  number  of  other,  not  very  well-defined  sulphides 
"with  more  or  less  complicated  formulae. 


CARBON  ;   OCCURRENCE.  269 


CHAPTER   XXXVII. 

CARBON. 

Carbon  ;  symbol,  C  ;  atomic  weight,  12 ;  specific  gravity,  as  diamond, 
3.5,  as  graphite,  2.14  to  2.35. 

THE  element  carbon  occurs  in  three  modifications,  two  of  which, 
diamond  and  graphite,  are  of  crystalline  structure,  while  the  third 
is  amorphous  carbon,  and  occurs  widely  distributed  in  the  form  of 
coal.  By  far  the  greater  quantity  of  carbon,  however,  is  found 
combined  in  the  numerous  compounds  of  that  element.  Carbon 
dioxide  (as  was  mentioned  on  page  167)  is  an  essential  constituent 
of  the  atmosphere.  The  carbonate  of  calcium  forms  limestone, 
chalk,  marble,  and  the  two  crystalline  minerals,  calcite  and  arragon- 
ite ;  the  combined  carbonates  of  calcium  and  magnesium,  under  the 
name  of  dolomite,  are  the  principal  structure  of  great  masses  of 
rock ;  and,  furthermore,  the  carbonates  of  iron,  zinc,  barium,  man- 
ganese, and  lead  are  important  additions  to  the  mineral  wealth  of 
the  world.  Carbon  is  also  invariably  present  in  all  of  the  innumer- 
able organic  compounds  with  which  we  are  acquainted ;  and,  further- 
more, the  products  of  vegetable  disintegration  which  are  classed 
under  the  head  of  coal  are,  in  the  main,  composed  of  the  element 
under  discussion. 

Carbon  is  dimorphous  ;  as  diamond  it  crystallizes  in  the  regular 
system,  while  as  graphite  it  is  monosymmetric.* 

The  greater  number  of  diamonds  occur  in  the  older  alluvial 
deposits,  but  some  have  been  found  imbedded  in  a  laminated  gran- 
ular quartz  rock  called  itacolumite  ;  they  are  also  sometimes  present 
in  a  species  of  conglomerate,  composed  of  rounded,  siliceous  pebbles, 
quartz,  and  chalcedony.  Diamonds  were  originally  imported  into 
Europe  from  the  East  Indies,  from  which  portion  of  the  world  and 
from  Borneo  the  only  specimens  were  procured  until  the  year  1727, 
when  large  diamond  fields  were  discovered  in  Brazil.  In  1867  the 

*  Formerly  supposed  to  be  hexagonal. 


270  DIAMOND;  OCCURRENCE,  PROPERTIES. 

diamond  fields  of  South  Africa  were  opened.  Some  diamonds  are 
also  found  in  the  Urals,  in  New  South  Wales,  and  in  the  United 
States. 

The  diamond  is  distinguished  by  its  extreme  hardness,  its 
great  power  of  refraction  and  brilliant  lustre ;  its  specific  gravity  is 
3.5  ;  it  is  a  poor  conductor  of  electricity  and  of  heat.  The  mineral 
also  occurs  in  black  pebbles  or  masses  known  as  carbonado,  having 
a  specific  gravity,  of  from  3.01  to  3.4.  A  coarse  variety  of  diamond 
which,  owing  to  imperfections  in  structure,  is  unfit  for  jewelry,  is 
sold  for  glass-cutting  purposes,  under  the  name  of  bort.  The  weight 
of  the  diamond  is  measured  in  carats  ( 1  carat  =  .205  gram) ;  the 
price  per  carat  increases  in  geometric  ratio,  although  always  modi- 
fied by  the  quality  of  the  stone.  The  largest  diamond  is  about  the 
size  of  half  a  hen's  egg ;  it  originally  weighed  800  carats,  but  was 
greatly  reduced  in  weight  by  cutting ;  the  Pitt  or  Regent  diamond 
weighs  136.75  carats,  and  is  of  unblemished  transparency  and  lustre. 
When  not  in  contact  with  the  air,  diamond  can  be  heated  to  a  white 
heat  without  alteration ;  when  heated  between  the  carbon  points  of 
an  arc  light,  it  swells  and  changes  to  a  grayish  mass  with  an  almost 
metallic  lustre  ;  in  this  form  it  resembles  ordinary  coke.  When 
heated  in  the  air,  diamond  takes  fire  at  about  1000°,  and  then  burns 
to  form  carbon  dioxide  (C02),  leaving  only  a  very  slight  trace  of 
ash.  Oxidizing  agents,  such  as  fused  potassium  nitrate,  or  potas- 
sium bichromate  and  sulphuric  acid,  can  oxidize  diamond  to  carbon 
dioxide.  Sir  Humphry  Davy  was  the  first  to  prove  that  diamond 
consisted  of  nearly  pure  carbon. 

Graphite,  also  called  plumbago  or  black  lead,  is  the  second  crys- 
talline form  of  carbon.  It  occurs  in  beds  and  imbedded  masses  in 
the  oldest  geologic  formations,  in  granite,  gneiss,  micaceous  schists, 
and  crystalline  limestone.  It  is  probably,  in  some  instances,  the 
resul^  of  the  alteration  of  deposits  of  coal  by  heat,  although  its 
origin  is  as  yet  imperfectly  understood.  In  some  places  the  graph- 
ite is  found  quite  pure  ;  for  instance,  in  the  "  Eureka  Black  Lead 
Mine  "  at  Sonora,  California,  there  is  a  bed  from  twenty  to  thirty  feet 
in  thickness,  which  contains  the  substance  in  so  pure  a  state  that  it 
can  be  cut  in  blocks  and  shipped  without  further  preparation.*  The 
ash  left  on  burning  this  graphite  is  only  about  five  per  cent  of  the 
whole.  Sometimes  the  graphite  is,  of  course,  much  more  impure, 
so  that  it  may  be  entirely  unfit  for  use.  The  chief  occurrences  of 
*  Practically  no  graphite  is  now  mined  in  California. 


GRAPHITE;  OCCURRENCE,  PROPERTIES.  271 

the  mineral  are  in  the  Urals  in  Siberia ;  in  Borrowdale,  Cumber- 
land ;  *  in  Arendal,  Norway ;  and  in  some  parts  of  Austria,  Russia, 
and  France ;  while  large  quantities  are  also  found  in  the  East  Indies. 
In  the  United  States,  the,  mineral  occurs  quite  frequently,  notably 
in  California ;  at  Sturbridge,  Mass.  ;  at  Ticonderoga,  and  in  the 
northern  part  of  Michigan.  Graphite  can  be  artificially  prepared 
by  crystallization  of  carbon  which  has  been  dissolved  in  melted 
iron ;  for,  when  gray  pig  iron  is  dissolved  in  acids,  the  insoluble 
graphite  remains  in  the  form  of  small,  delicate  scales  ;  a  similar 
form  of  the  substance  has  also  been  discovered  in  some  meteorites. 

Graphite  is  used  in  the  manufacture  of  lead  pencils,  infusible 
crucibles  and  as  a  lubricator ;  it  is  adapted  to  the  latter  purpose 
because  the  substance  is  soft  and  scaly.  It  is  grayish  black,  and 
has  almost  a  metallic  lustre.  When  burned  in  the  air,  it  forms  car- 
bon dioxide,  and  leaves  an  ash,  which,  when  it  is  derived  from  the 
purer  varieties,  consists  mainly  of  ferric  oxide  and  silica.  When 
heated  with  concentrated  nitric  acid  for  some  length  of  time,  graph- 
ite changes  to  a  yellow,  crystalline  body,  which  contains  carbon, 
hydrogen,  and  oxygen.  This  substance  is  known  as  graphitic  acid. 
Graphitic  acid,  when  heated,  disintegrates  almost  with  explosive 
violence,  leaving  a  voluminous  black  residuum,  which  apparently 
consists  of  very  finely  divided  graphite.  This  latter  form  of  the 
substance  is  applied  as  a  covering  to  the  moulds  used  in  electroplat- 
ing ;  for,  as  graphite  is  a  good  conductor  of  electricity,  it  renders 
the  surfaces  of  the  non-conducting  substances  from  which  these  are 
made,  capable  of  conducting  electricity,  the  conductivity  of  the  sur- 
face being  an  essential  preliminary  to  forming  a  metallic  deposit. 

Those  compounds  which  are  formed  in  animal  and  vegetable 
organisms,  and  which  are  classed  under  the  general  head  of  organic 
substances,  are  produced  by  the  union  of  a  very  few  elements ; 
namely,  carbon,  hydrogen,  oxygen,  nitrogen,  sulphur,  and  pjjos- 
phorus.  When  such  substances  decompose"  in  the  open  air  they 
break  down  completely,  changing  for  the  most  part  into  gaseous 
products  ;  but  when  the  vegetable  fibres  are  protected  by  a  layer  of 
water,  as  is  the  case  in  peat-bogs,  the  process  of  decomposition  goes 
on  slowly ;  certain  portions  of  the  constituents  of  the  organic  sub- 
stances, especially  oxygen  and  hydrogen,  generally  pass  off  in  other 
combinations,  while  the  vegetable  substance  becomes  changed,  first 
into  peat  and  then  into  bituminous  coal,  and,  at  the  same  time,  the 
*  The  mines  of  Borrowdale  are  now  exhausted. 


272 


COAL;    FORMATION. 


percentage  of  contained  carbon  increases.  Peat,  brown  coal,  bitu- 
minous coal,  and  anthracite  coal  are  successive  steps  in  the  process 
of  floral  decomposition ;  when  the  anthracitic  stage  is  reached  the 
changes  have  become  so  complete  that  a  black,  shiny,  homogeneous 
mass  has  resulted ;  in  this  mass  the  original  vegetable  structure  has 
entirely  disappeared,  or  is,  at  least,  so  indistinct  that  special  means 
must  be  taken  for  its  detection.  The  pressure  to  which  the  dead 
organic  structures  are  subjected  is  of  material  influence  on  the 
rapidity  with  which  a  peat  formation  is  changed  to  anthracite  ;  in- 
deed, in  districts  of  Russia  where  the  coal  has  not  been  placed 
under  very  great  pressure,  a  brown  coal  (lignite),  which  can 
scarcely  be  distinguished  from  peat,  is  found  in  places  where  the 
age  of  the  deposit  would  lead  one  to  expect  anthracite.  A  compari- 
son of  the  approximate  composition  of  the  combustible  portions 
of  some  of  the  varieties  of  coal  will  show  the  changes  which 
the  vegetable  matter  undergoes  during  its  decomposition  more 
clearly  :  — 


CARBON,  PEE  CENT. 

HYDROGEN,  PER  CENT. 

OXYGEN  ANU  NITROGEN, 
PER  CENT. 

Wood 

50 

6 

44 

Peat 

60 

5.75 

34.25 

Lignite 
Bituminous  coal 

67 

87 

5.3 

5.6 

27.7 
7.4 

Anthracite 

94 

3.4 

2.6 

The  various  forms  of  coal  are  amorphous,  and  therefore  differ 
markedly  from  diamond  or  graphite,  both  of  which  are  crystalline. 

When  organic  substances  are  heated  without  access  of  air  they 
undergo  a  process  of  carbonization;  the  volatile  products  of  this 
destructive  distillation  pass  off  as  gases  and  liquids  (see  page  183), 
while  amorphous  carbon  is  left  behind.  Similar  changes  take  place 
during  the  destructive  distillation  of  bituminous  coal,  leaving  coke.** 
Coke  is  a  porous,  shiny  form  of  amorphous  carbon ;  it  conducts  elec- 
tricity and  heat  about  as  well  as  graphite.  Coke  may  contain  as 
much  as  91.5  per  cent  of  its  total  weight  in  the  form  of  carbon. 

Gas  coal,  which  collects  on  the  walls  of  retorts  in  which  bitu- 
minous coal  is  heated  to  form  illuminating  gas,  is  a  product  of  the 
decomposition  of  gaseous  compounds  of  carbon  and  hydrogen.  It 


CARBON;   AMORPHOUS.  273 

has  almost  a  metallic  lustre,  resembling  very  dense  coke ;  it  is  diffi- 
cult to  ignite,  and  conducts  heat  and  electricity  quite  well. 

Wood  charcoal  is  produced  by  the  imperfect  combustion  of  wood 
sticks,  animal  charcoal  by  a  similar  treatment  of  animal  refuse, 
such  as  bones  or  blood.  The  finer  forms  of  bone  charcoal  are 
termed  bone  black  and  ivory  black.  All  forms  of  charcoal,  but 
especially  the  varieties  of  animal  charcoal,  have  a  remarkably  pro- 
nounced tendency  to  absorb  coloring  matters  from  solutions.67 
These  colored  substances  are  apparently  deposited  within  the 
porous  substance  of  the  coal,  for  they  can  be  extracted  unchanged 
therefrom  by  means  of  the  proper  solvents.*  The  property  of 
absorbing  coloring  matter  does  not  belong  to  charcoal  alone ;  all 
insoluble  porous  substances  can  perform  the  same  office  in  a  greater 
or  less  degree ;  f  in  rare  instances  it  may  happen  that  the  charcoal 
exercises  a  reducing  action  on  the  absorbed  matter.  Crude  sugar  is 
decolorized  by  means  of  charcoal. 

The  purest  form  of  amorphous  carbon  is  lamp-black,  which 
results  from  the  combustion  of  carbon  and  hydrogen  compounds 
where  an  imperfect  supply  of  oxygen  is  provided,  or  where '  the 
flame  is  cooled  before  perfect  combustion  has  taken  place ;  lamp- 
black is  therefore  deposited  on  a  cold  porcelain  or  metal  plate 
placed  within  a  luminous  gas  or  lamp  flame.  The  lamp-black  of 
commerce  is  obtained  by  burning  resinous  pine  wood,  tar,  or  some 
kinds  of  bituminous  coal.  The  substance  is  collected  on  coarse 
cloths  hung  over  the  burning  wood  placed  in  suitable  chambers. 
Lamp-black  is  used  in  the  manufacture  of  printers'  and  Indian  ink. 

*  Indigo,  which  has  been  dissolved  in  sulphuric  acid  and  absorbed  from 
this  solution  by  charcoal,  can  be  extracted  from  the  charcoal  by  alkalies.  Me- 
tallic oxides,  absorbed  by  charcoal,  can  be  extracted  by  strong  acids. 

t  Aluminium  hydroxide,  ferric  hydroxide,  or  precipitated  sulphide  of  lead 
can  absorb  coloring  matter. 


274  METHANE;   OCCURRENCE. 


CHAPTER   XXXVIII. 

THE  COMPOUNDS  OF  CARBON  WITH  HYDROGEN. 

Methane  ;  formula,  CH4;  specific  gravity,  air  =  1,  is  .5531,  H2=  2, 
is  15.93  ;  molecular  weight,  16.032.  1  c.c.  of  the  gas  at  0°  and 
760  m.m.  weighs  .0007153  gram. 

THE  simplest  hydrogen  compounds  of  the  carbon  family  have 
the  formula  XH4,  where  X  represents  an  atom  of  some  element  of 
that  family;  as  a  consequence  the  valence  toward  hydrogen  pos- 
sessed by  the  atoms  of  the  elements  of  this  group  is  greater  by  one 
than  is  the  valence  of  those  in  the  preceding  (nitrogen)  family. 

Methane,  or  marsh  gas  (the  hydrogen  compound  of  carbon  which 
corresponds  to  ammonia  in  the  nitrogen  family),  occurs  quite  fre- 
quently in  nature  as  a  product  of  the  decay  of  vegetable  tissues. 
The  muddy  bottom  of  any  stagnant,  marshy  pool,  when  stirred, 
emits  bubbles  of  marsh  gas,  which,  however,  always  contain  from 
ten  to  twenty  per  cent  of  carbon  dioxide,  as  well  as  a  small  amount 
of  nitrogen.  The  metamorphoses  which  resulted  in  the  formation 
of  coal  beds,  having  been  similar  to  the  changes  taking  place  in 
marshes,  must  necessarily  also  have  produced  methane,  so  that,  as  a 
consequence,  pockets  of  the  gas  (sometimes  subjected  to  great  pres- 
sure) not  infrequently  occur  in  coal  mines  ;  the  escaping  gas,  when 
a  pocket  is  tapped,  forms  a  dangerously  explosive  mixture  with  the 
air.  Methane,  in  some  places,  escapes  from  openings  in  the  ground  ; 
the  gas  which  is  passing  off  is  sometimes  either  intentionally  or 
accidentally  ignited  ;  the  burning  gas  wells  so  produced  are,  in  some 
instances,  regarded  with  superstitious  reverence,  as  is  the  case  with 
the  holy  fire  at  Baku  on  the  Caspian  Sea. 

The  natural  gas  which  is  used  so  extensively  for  illuminating 
and  heating  purposes  in  a  number  of  places  in  the  United  States 
consists,  for  the  most  part,  of  methane.  Methane  is  always  pro- 
duced in  large  quantity  by  the  dry  distillation  of  coal  ;  it  therefore 
forms  the  major  portion  of  illuminating  gas. 


METHANE;    PREPARATION.  275 

Methane  cannot  be  prepared  by  direct  union  of  the  elements 
carbon  and  hydrogen,  yet  if  hydrogen  can  be  brought  to  act  upon 
carbon  when  the  latter  is  in  what  may  be  considered  the  nascent 
state,  then  the  circumstances  are  such  that  union  can  take  place. 
An  example  of  such  a  production  of  methane  is  found  in  the  reac- 
tion between  carbon  monoxide  (CO)  and  hydrogen,  under  the  in- 
fluence of  a  strong  discharge  of  electricity  from  an  induction  coil  :  — 


The  Preparation  of  Methane. 

The  best  method  to  prepare  methane  for  laboratory  use  is  by  the 
dry  distillation  of  some  organic  substance,  and,  as  we  have  seen, 
our  choice  of  these  is  not  very  limited.  Experience  has  shown, 
however,  that  a  mixture  of  some  dry  acetate  with  a  strong  base 
will  yield  methane  ;  an  example  of  the  production  of  the  gas  in  this 
way  can  be  found  in  the  decomposition  of  the  acetate  of  sodium,  when 
that  substance  is  heated  in  the  presence  of  sodium  hydroxide  :  — 

CH.COONa+NaOH  =  CH4     +Na2C03 

Sodium  acetate  +  Sodiurn  hydroxide  =  Methane  +  Sodium  carbonate. 
Sodium  acetate  can  be  considered  as  methane  in  which  one  hydro- 
gen atom  has  been  replaced  by  the  univalent  group  —  COONa. 

Methane  is  a  colorless  gas,  without  odor  or  taste.  Its  specific 
gravity  is  .5531.  Under  a  pressure  of  one  atmosphere,  methane 
boils  at  —  164°  ;  if  evaporated  quickly  under  diminished  pressure 
the  liquid  will  be  cooled  to  below  its  freezing  point,  and  will  form 
a  snow-like  mass.  The  gas  burns  in  oxygen  or  air  with  a  nearly 
colorless  flame  which  is  much  like  that  of  hydrogen,  while  carbon 
dioxide  and  water  are  formed  :  — 


Methane  is  decomposed  into  its  constituent  elements  only  at 
quite  a  high  heat  ;  when  passed  through  a  white-hot  tube  it  breaks 
down  into  carbon  and  hydrogen  ;  it  is,  therefore  (with  the  exception 

*  A  similar  and  most  interesting  production  of  methane  is  by  the  action  of 
copper  on  a  mixture  of  the  vapors  of  carbon  bisulphide  and  hydrogen  sulphide  :  — 

8Cu  +  2H2S  +  CS2=4Cu2S  +  CH4. 

Here  the  carbon  and  the  hydrogen  may  both  be  considered  to  be  in  the  nascent 
state  ;  the  copper  simply  removes  sulphur  and  leaves  the  carbon  and  hydrogen 
to  rearrange  themselves  into  the  most  stable  configuration  under  existing  con- 
ditions. (See  page  51.) 


276 


METHANE  ;    VOLUMETEIC    COMPOSITION. 


of  ammonia),  more  stable  than  any  of  the  hydrogen  compounds  of 
the  preceding  family.  Methane,  if  it  is  mixed  with  exactly  enough 
oxygen  to  form  carbon  dioxide,*  can  be  exploded  by  means  of  an 
electric  spark ;  if  care  is  taken  to  keep  the  water  vapor  produced 
by  this  reaction  in  the  form  of  a  gas,  the  result  is  that  one  volume 
of  methane  with  two  of  oxygen  forms  one  volume  of  carbon  dioxide 
and  two  of  water;  the  total  volume  of  the  mixture  of  gases  is 
therefore  the  same  after  the  explosion  as  it  was  before :  — 


1  vol.  methane,  +      2  vols.  oxygen,        =  1  vol.  carbon  dioxide +  2  vols.  water  vapor. 

From  these  results  it  follows  that  one  molecule  of  methane  is 
able  to  form  two  molecules  of  water  vapor,  and,  consequently,  as 
two  molecules  of  water  vapor  contain  four  atoms  of  hydrogen, 
methane  must  also  have  four  hydrogen  atoms  in  its  molecule.  The 
specific  gravity  of  methane,  hydrogen  =  2,  is  15.9;  this  shows  that 
its  corrected  molecular  weight  must  be  16,  f  for  when  methane  is 
analyzed  we  find  it  composed  of  12  parts  by  weight  of  carbon  and  4 
parts  of  hydrogen,  which  form  16  parts  of  methane.  The  4  parts  of 
hydrogen,  as  we  have  seen,  represent  four  atoms;  that  12  parts  of 
carbon  represent  one  atom  we  presume  to  be  the  case  because,  in  no 
compound  of  carbon,  the  specific  gravity,  and  hence  the  molecular 
weight,  of  which  is  known,  has  carbon  ever  been  found  to  enter  with 
a  less  proportional  weight  than  twelve.  After  considering  these 
experimental  facts,  we  conclude  that  the  formula  of  methane  is  CH4. 
(  See  page  187.) 

When  methane  is  mixed  with  chlorine  and  placed  in  the  dark, 
no  reaction  takes  place,  but  when  the  mingled  gases  are  exposed  to 
the  sunlight,  a  violent  explosion  results ;  hydrochloric  acid  and  car- 
bon being  produced.  This  action  is  exactly  parallel  to  the  action 
of  chlorine  on  all  other  hydrogen  compounds  (with  the  exception  of 
hydrofluoric  and  hydrochloric  acids)  :  — 

*  A  eudiometer  tube  is  used  for  this  purpose.  See  Note  20,  laboratory 
appendix. 

t  In  exact  numbers  16.032,  because  the  atomic  weight  of  hydrogen  is 
1.008;  in  considerations  of  this  kind  it  is  better  to  neglect  the  decimal. 


METHANE;   SUBSTITUTION.  277 

CH4  +  4  Cl  =  C  +  4  H  Cl, 
NH3  +  3Cl"=N  +3HC1, 
SH2  +  2  Cl  =  S  +  2  H  Cl, 
BrH  +  Cl  =  Br  +  H  Cl. 

When,  however,  the  chlorine  is  allowed  to  attack  methane 
slowly,  as  it  does  in  diffused  light,  substitution  of  hydrogen  for 
chlorine  results,  so  that  the  following  changes  successively  take 
place :  — 

CH4  +  2  Cl  =  CH3  Cl  +     HC1, 

CH4  +  4  Cl  =  CH2  C12  +  2  H  Cl, 

CH4  +  6  Cl  =  CH  Cls  +  3  H  Cl, 

CH4  +  8C1=      CC14  +  4HC1. 

When  one  hydrogen  atom  has  been  removed  from  methane,  the  un- 
saturated  univalent  radicle  (see  page  108)  is  termed  methyl,  when 
two  hydrogen  atoms  are  removed  the  bivalent  radicle  is  methylen, 
and  when  three  hydrogen  atoms  are  removed  the  trivalent  radicle  is 
methin : 

CH3— ,  methyl;  CH3C1,  methyl  chloride. 

CH2  =,  methylen ;  CH2  C12 ,  methyleii  chloride. 

CH   ==,  methin;  CHC13,  methin  chloride  (chloroform). 

These  chlorinated  substances  can,  therefore,  all  be  considered  as 
methane  in  which  one,  two,  or  three  atoms  of  hydrogen  have  re- 
spectively been  replaced  by  chlorine ;  they  partake  more  or  less  of 
the  nature  of  methane,  although  the  introduction  of  successive 
chlorine  atoms  causes  the  resulting  compound  to  depart  more  and 
more  from  the  character  of  the  type,  thus :  — 

Methyl  chloride  is  a  gas,  colorless,  becomes  liquid  at  —  23°. 7. 
Methylen  chloride,  liquid,  boils  at  40°. 
Methin  chloride,  liquid,  boils  at  61°.2. 
Carbon  tetrachloride,  liquid,  boils  at  76°.5. 

With  the  introduction  of  each  chlorine  atom  the  boiling  point  in- 
creases, and,  therefore,  each  of  these  changes  brings  the  character  of 
the  chlorine  substituted  methane  farther  from  that  of  the  colorless 
gas  from  which  it  is  derived. 

When  hydrogen  is  removed  from  methane,  the  resulting  unsat- 
urated  monovalent  group,  methyl,  cannot  exist  alone  ;  in  this  respect 
it  is  like  a  free  atom  of  hydrogen,  and,  therefore,  methyl  seeks  the 
first  opportunity  of  uniting  with  some  atom  or  radicle  to  form  a 


278  HIGHER    HYDROCARBONS. 

saturated  compound.  We  have  seen  that,  when  chlorine  is  present, 
the  methyl  reacts  with  that  element  to  form  methyl  chloride ;  if  no 
such  other  substance  with  which  methyl  is  capable  of  union  can  be 
found,  the  radicle  will  join  with  itself  to  form  dimethyl  (ethane)  ; 
CH3  —  CH3.*  Dimethyl,  like  methane,  is  capable  of  having  from 
one  to  six  of  its  hydrogen  atoms  substituted  by  chlorine.  The  first 
reaction  which  takes  place  between  ethane  and  chlorine  is  as 
follows :  — 

CH3—  CH3  +  2  Cl  =  CH3  —  CH2  Cl  +  H  Cl. 

The  compound  CH3  —  CH2C1  (C2H5C1)  is  termed  ethyl  chloride 
(the  same  system  of  nomenclature  applies  in  this  case  as  it  does 
with  the  methyl  compounds,).  If  we  remove  one  hydrogen  atom  from 
ethane  to  form  the  monovalent,  unsaturated  radicle  ethyl,  the  latter, 
if  no  other  substance  is  present  with  which  it  can  unite,  will  form 
diethyl  ( butane)  just  as  methyl  forms  dimethyl  (ethane)  :  — 

CH3  —  CII3  -  H  =  CII3  —  CH2 ;  CH3  —  CH2  —  +  CII3  —  CH2  —  =  CH3  —  CHa  —  CHa  —  CH3. 

Ethane— Hydrogen  =  ethyl.          Ethyl  +  ethyl        =  diethyl  (butane). 

Diethyl  can  likewise  have  its  hydrogen  atoms  substituted  by  chlo- 
rine, and,  by  the  loss  of  one  atom  of  hydrogen,  can  be  converted  into 
the  monovalent,  unsaturated  radicle  (butyl),  which  further  unites 
with  itself  to  form  dibutyl  or  octane.  It  is  possible,  however,  so  to 
modify  the  above  reaction  as  to  bring  ethyl  and  methyl  together,  in 
which  case  the  two  radicles  will  unite  to  form  ethyl-methyl  (pro- 
pane) t ;  in  the  same  way,  a  mixture  of  propyl  and  ethyl  will  yield 
propylethyl  (pentane).  By  a  judicious  combination  of  the  iodides 
of  organic  radicles,  carbon  and  hydrogen  compounds  containing  as 
many  as  sixty  carbon  atoms  in  a  molecule  t  have  been  prepared.  These 
substances,  as  they  contain  only  carbon  and  hydrogen,  are  termed 

*  Methyl  iodide  boiled  with  zinc  dust  or  with  sodium  forms  ethane.     The 
reaction  takes  place  as  follows :  — 

CH3I  +  Zn  +  CH3I=ZnI2  +  CH3  +  CH3.     CH8  +  CH3  =CH3— CH3. 
This  reaction  is   also  applicable  in  the  formation  of  the  more  complicated 
compounds  which  follow.    Of  course,  did  we  wish  to  prepare  some  of  the  latter, 
we  would  not  use  methyl  iodide  and  zinc,  but  would  employ  the  iodides  of  those 
radicles  which  we  wish  to  unite. 

t  By  boiling  a  mixture  of  ethyliodide  and  methyliodide  with  zinc-dust 
thus  :  — 

CH3  —  CH2I  +  Zn    +CH3I  =  ZnI2  +  CH3  — CH2  — CH3. 

Ethyliodide  +  methyliodide  =  zinc  iodide  +  ethylmethyl  (propane). 

J  Hell  and  Hagele ;  Berichte  der  Deutsch.  Chem.  Gesell. ;  22,  502.     This 
hydrocarbon  (Cm  H122)  was  produced  by  heating  C30  H61 1  with  sodium. 


PETllOLEUM. 


, 

hydrocarbons,  the  particular  class  of  saturated  hydrocarbons  now 
under  discussion  being  called  paraffins.  The  first  seven  represen- 
tatives of  this  class  are  given  in  the  following  table :  — 


BOILING  POINT. 

SPECIFIC  GRAVITY 

OF  LIQUID. 

Methane,  CH4 

-164° 

0.415  (at  - 

164°) 

Ethane,  C.2  H6 

Propane,  L/3  ±18 
Butane,  C4H10 

+  1° 

0.6 

Pentane,  C5H12 

+  37° 

0.627 

Hexane,  C6H14 

+  69° 

0.658 

Heptane,  C7H16 

+  98° 

0.683 

A  general  formula  for  these  compounds  is  CnH2n  +  2,  where  n  rep- 
resents the  number  of  carbon  atoms  in  the  chain  ;  an  increase  of  this 
number  by  one  in  any  given  paraffin  of  the  series  under  discussion 
raises  the  boiling  point  of  that  paraffin  by  about  19°.  The  com- 
pounds which  begin  the  series  are  gases,  those  with  from  five  to 
sixteen  carbon  atoms  are  liquids  at  ordinary  temperatures,  the  re- 
mainder are  solids  with  melting  points  ranging  from  18°  to  74°,  * 
the  specific  gravities  of  the  hydrocarbons  increase  with  the  number 
of  carbon  atoms  in  the  molecule,  but  are  always  less  than  unity. 

The  hydrocarbons  Cn  H2  w  +  2  are  found  in  coal  oil ;  the  latter  is 
generally  technically  divided  as  follows  :  — 

Petroleum  ether,  boiling  point  50°  to   70°;  pentane  and  hexane. 
Benzine,  u  70°  to   90°;  hexane  and  heptane. 

Ligroine,  "  90°  to  120°;  heptane  and  octane. 

Petroleum,  (kerosene)  "         150°  to300°;  octane  to  hexadecane  (C^H^). 
The  higher  boiling  portions  are  vaseline  and  paraffin. 

If  one  hydrogen  atom  is  removed  from  ethane,  there  results  an 
un saturated  radicle  (ethyl)  which  cannot  exist  alone  ;  we  saw  that  it 
unites  with  itself  to  form  diethyl  (  butane).  Experience  has  shown 
us,  however,  that  an  entirely  different  result  may  be  expected  if  we 
simultaneously  remove  a  hydrogen  atom  from  each  of  the  carbon 
atoms  contained  in  ethane ;  the  molecule  containing  the  pair  of 
neighboring  carbon  atoms,  which  have  thus  become  unsaturated,  is 
then  capable  of  independent  existence  and  is  called  ethylene :  — 

*  A  few  compounds  recently  discovered  may  prove  to  be  an  exception  to 
this  rule.  The  hydrocarbon  with  sixty  atoms  of  carbon  in  one  molecule  melts 
at  102°. 


280  ETHYLENE. 

CH3  —  CH3,  ethane,  CH2  —  CH2,  ethylene. 

What  is  true  of  ethylene  remains  true  when  the  hydrogen  atoms 
of  that  substance  are  substituted  by  other  atoms  or  groups  of  atoms  ; 
we  can,  therefore,  beginning  with  ethylene  as  a  nucleus,  by  repla- 
cing the  hydrogen  atoms  with  ethyl,  methyl  propyl,  etc.,  construct  a 
new  series  of  unsaturated  carbon  compounds  which  would  have  the 
general  formula  Cn  H2  n  .  A  further  discussion  of  these  complicated 
substances  belongs  in  the  domain  of  organic  chemistry.  The  fact 
that  carbon  atoms  are  never  known  to  be  unsaturated  in  organic 
compounds  unless  the  unsaturated  atoms  are  side  by  side  *  and  the 
consideration  that  carbon  is  tetravalent  in  methane,  have  led  chem- 
ists to  regard  the  carbon  atoms  in  ethylene  as  being  joined  to  each 
other  in  a  different  way  from  that  in  which  they  are  in  ethane  ;  for, 
if  we  suppose  the  carbon  atoms  to  be  always  quadrivalent,  then  the 
pair  of  carbon  atoms  in  ethane  are  joined  by  one  valence  of  each 
atom,  where  in  ethylene  they  are  united  by  two.  The  following 
diagram  will  make  this  more  clear  :  — 
H  H 

H  H 


H  —  C  —  C  —  H  cCx 

I         I  K'  XH; 

H      H 
Ethane.  Ethylene. 

These  structural  formulae  express  the  theory  that  each  valence  of 
any  carbon  atom  (which  is  of  necessity  tetravalent)  must  neutralize 
a  corresponding  valence  of  some  other  atom.  All  we  can  really 
know  regarding  these  combinations  is  that  the  carbon  atoms  in  any 
such  compound  as  ethylene  are  held  together  by  a  certain  force,  of 
the  nature  of  which  we  are  ignorant,  and  which  we  call  chemism  or 
chemical  affinity.  It  is  usually  stated  that  "  carbon  has  four  points 
of  affinity,  or  four  valences;"  of  course,  provided  we  consider 
chemism  as  a  force,  such  a  theory  is  not  tenable,  because  no  force 
can  act  unless  it  has  something  to  act  upon  ;  when  the  carbon  atoms 
are  united  as  they  are  in  ethylene  a  certain  amount  of  residual  force 

*  A  few  compounds  in  organic  chemistry  admit  of  a  different  interpreta- 
tion for  their  structural  formulae  ;  it  seems  not  improbable,  for  example,  that 
a  hydrocarbon  of  the  formula  —  CH2  —  CH2  —  CH2  —  is  capable  of  existence. 
In  such  a  compound  the  two  carbon  atoms  at  the  ends  of  the  chain  would  be 
unsaturated  and  trivalent. 


ACETYLENE.  281 

(beyond  that  required  to  hold  these  two  atoms  together)  acts  upon 
and  retains  the  four  hydrogen  atoms  ;  this  much  we  know  from 
experimental  evidence ;  but  if  we  further  suppose  that  the  carbon 
atoms  are  each  joined  by  two  points  of  affinity,  we  must  then  ac- 
cept the  proposition  that  the  force  which  unites  the  two  atoms  is 
manifest  only  from  four  distinct  spots  upon  their  surfaces,  an  hy- 
pothesis which  is  not  in  accord  with  what  we  know  as  regards  the 
attraction  which  one  mass  exercises  toward  another.  A  statement 
which  would  more  nearly  accord  with  our  experimental  knowledge 
would  be  that  in  ethylene  we  have  two  trivalent,  and  hence  unsat- 
urated,  carbon  atoms  joined  to  each  other.*  If  by  some  means  we 
remove  a  hydrogen  atom  from  each  of  the  neighboring  carbon  atoms 
in  ethylene,  there  results  a  compound  CH  —  CH=(C2H2)  which 
is  termed  acetylene.  The  carbon  atoms  in  acetylene  are  supposed 
to  be  united  by  a  so-called  "  triple  linking,"  for  the  theory  which 
was  used  in  explaining  the  constitution  of  ethylene  must  compel  the 
supposition  that  the  two  additional  unsaturated  valences  in  acetylene 
must  neutralize  each  other ;  the  formula  of  the  latter  substance  is, 
therefore,  written  CH  =  CH.  In  this  case  we  also  must  indulge  in 
speculation  if  we  wish  to  go  farther  than  to  assume  any  more  than 
the  existence  of  two  divalent  unsaturated  carbon  atoms  in  acetylene. 
Acetylene  can  act  as  the  nucleus  of  another  series  of  hydrocarbons 
with  the  general  formula  CnH2n_2,  for,  by  substituting  the  hydro- 
gen atoms  in  acetylene  by  organic  radicles  (methyl,  ethyl,  etc.),  we 
can  produce  long  chains  of  atoms.  The  hydrocarbons  of  the  series 
CuH2n  and  Cn  H2n_2,  being  unsaturated,  can  readily  add  substances 
such  as  chlorine,  bromine,  hydrobromic  acid,  etc. 

Ethylene  is  a  colorless  gas  which  is  poisonous ;  it  has  a  specific 
gravity  of  .9852,  is  tolerably  soluble  in  water,  and  when  heated 
breaks  down  into  methane  and  acetylene  :  — 

*  Recent  investigations  in  organic  chemistry  seem  to  show  that  carbon 
atoms  react  as  if  their  attractive  force  were  exerted  along  four  lines  con- 
necting the  centre  of  a  sphere  with  four  points  symmetrically  grouped  upon 
its  surface;  these  four  points  would  then  correspond  to  the  angles  of  a  regular 
tetrahedron.  The  hydrogen  atoms  in  methane  would  then  be  at  the  points  of 
the  tetrahedron,  and  the  two  carbon  atoms  in  ethylene  would  be  joined  along 
the  line  connecting  two  of  these  points.  By  this  hypothesis  the  difference 
between  saturated  and  unsaturated  carbon  chains  can  be  explained.  See  also 
M.  M.  Pattison  Muir;  Principles  of  Chemistry.  The  discussion  here  given 
has  been  set  forth  in  detail  by  Lossen;  Liebig's  Annalen;  204,  265. 


282  THE   FLAME. 


Owing  to  this  latter  decomposition  ethylene  burns  with  a  lumi- 
nous flame,  because  acetylene  is  a  gas  which  undergoes  decomposition 
into  methane  and  carbon  :  — 

2C2H2  =  CH4  +  3C, 

and  therefore  it  emits  a  luminous  flame,  for  in  each  case  the  glow- 
ing particles  of  carbon  emit  the  light. 

Illuminating  gas,  prepared  by  the  distillation  of  soft  coal,  is 
composed  chiefly  of  hydrogen,  methane,  ethylene  and  acetylene, 
carbon  monoxide  and  carbon  dioxide,  and  nitrogen.  The  quality  of 
the  flame  is  determined  by  the  amount  of  ethylene  and  acetylene 
present,  for  a  gas  which  contains  these  unsaturated  hydrocarbons 
burns  with  a  luminous  flame,  while  methane  or  carbon  monoxide 
scarcely  gives  any  light  during  combustion. 

A  flame  can  be  observed  wherever  a  gas,  in  consequence  of 
chemical  action,  is  heated  sufficiently  to  cause  it  to  glow.  In  most 
cases  this  chemical  reaction  is  caused  by  chemical  union  ;  that  it  is 
possible,  however,  to  have  a  flame  resulting  from  the  heat  of  decom- 
position of  an  endothermic  compound  is  proved  by  the  appearance 
of  a  flash  accompanying  the  explosion  of  nitrogen  chloride.  The 
simplest  case  of  flame  production  can  be  illustrated  by  mixing  two 
gases,  which  are  capable  of  giving  off  a  large  amount  of  heat  in 
their  union,  and  igniting  the  mixture  by  means  of  an  electric  spark 
or  by  means  of  a  flame.  In  this  mixture  the  molecules  of  the  two 
gases  are  intimately  intermingled,  the  reaction  takes  place  almost 
simultaneously  at  all  points  throughout  the  volume  of  the  gas,  and 
is,  therefore,  accompanied  by  an  explosion  and  the  formation  of  a 
homogeneous  flame.  f 

The  most  common  form  of  flame  is  produced  by  a  stream  of  gas 
pouring  into  a  volume  of  another  with  which  it  can  chemically 
unite  ;  if  the  entering  gas  is  heated  to  its  kindling  temperature, 
then  union  will  take  place  along  the  boundary  where  the  two  gases 
touch.  The  conical  shape  assumed  by  the  flame  of  a  gas  escaping 
from  a  round  vent  is  caused  by  the  diminution  in  the  quantity  of 
that  gas  as  the  distance  from  the  opening  increases,  this  diminution 
being  caused  by  the  consumption  of  the  gas  in  burning.  The 
structure  of  the  flame  of  a  gas  burning  in  air  can  be  taken  as  a 
type  of  all  others.  Such  a  flame  consists  of  a  number  of  zones 
which  can  be  easily  distinguished  by  their  appearance.  The  flame 


THE   FLAME. 


283 


of  a  candle,  for  instance,  exhibits  a  dark  centre  which  is  surrounded 

by  a  conical,  luminous  zone  ;  a  piece  of  paper  placed  over  this,  as 

is  shown  in  Fig.  11,  will  be  charred  in  the  form  of  a  circle,  this 

experiment  showing  that  the  gases  in 

the  centre  of  the  flame  are  not  heated 

to  a  high  temperature.    It  is  also  true 

that   a   small    piece   of   phosphorus 

placed  in  the  dark  centre  will  not 

burn  ;  so  there  can  be  no  oxygen  pres- 

ent.    Outside  of   this   central   cone 

there   is   a   luminous    zone  of  some 

thickness  where  the  oxygen  is  unit- 

ing with  the  escaping  gases  ;  but,  as 

oxygen  cannot  enter  into  this   por-  Fig.  n. 

tion  of  the  flame  in  excess,  the  carbon,  separated  from  the  glowing 

gases*  by  reason  of  the  high  temperature  of  the  flame,  is  not  com- 

pletely burned,  but  is  only  heated  to  a  white  heat.     Oxygen  is  pres- 

ent in  excess  on  the  outer  surface  of  the  luminous  zone,  and  therefore 

combustion  is  most  energetic  in  this  division  of  the  flame,  so  that 

an  enveloping  mantle,  which  is  scarcely  visible,  results  ;  this  is  the 

hottest  part  of  the  flame.      The  glowing  carbon  which  causes  a 

flame  to  become  luminous  is  produced  by  the  breaking  down  of 

the  heated  hydrocarbons  present  in  the   gas,  a  change  similar  to 

that  undergone  by  acetylene  taking  place  :  — 


From  this  it  follows  that  if  the  temperature  of  the  flame  can  be 
lowered  to  such  a  point  as  to  prevent  this  decomposition,  the  flame 
will  become  non-luminous  ;  such  an  alteration  can  be  brought  about 
by  diluting  the  gases  (before  burning)69  with  some  indifferent  sub- 
stance, such  as  carbon  dioxide,  and  the  same  result  can  be  accom- 
plished by  providing  a  supply  of  air  to  the  illuminating  gas  before 
the  vent  at  which  the  flame  is  lighted  is  reached.  The  Bunsen 
burner  attains  this  end  by  causing  the  gas  which  is  escaping  from 
a  small  central  opening  to  traverse  a  wider  brass  tube  before  igni- 

*  See  ethylene  and  acetylene.  It  seems  probable  that  ethylene  breaks  down 
into  acetylene  and  methane,  that  acetylene  condenses  to  form  more  complicated 
hydrocarbons,  that  the  latter  finally  yield  carbon  and  hydrogen.  Methane,  at 
high  temperatures,  also  yields  carbon  and  hydrogen. 


284  THE   FLAME. 

tion.  At  the  bottom  of  this  brass  tube  two  holes  are  pierced,  allow- 
ing the  entrance  of  a  limited  supply  of  air.  This  air  mingles  with 
the  escaping  gases,  and  thus  provides  for  complete  combustion 
before  the  decomposition  of  the  hydrocarbons  takes  place.  Un- 
doubtedly the  non-luminous  character  of  the  Bunsen  flame  is  also, 
in  part,  brought  about  by  the  addition  of  the  indifferent  gas  nitro- 
gen, which  must  necessarily  enter  the  burner  in  company  with 
oxygen.  If  the  gases  composing  a  flame  can  be  cooled  below  their 
kindling  temperature,  the  flame  will  be  extinguished. 

From  what  has  been  said  regarding  the  formation  of  a  flame  it 
follows  that  it  is  a  matter  of  indifference  which  of  the  two  gases 
uniting  to  form  the  flame  is  entering,  and  which  forms  the  surround- 
ing medium,  for  the  phenomena  are  caused  solely  by  the  union 
of  the  two.  The  terms  "  combustible "  and  "  a  supporter  of  com- 
bustion," as  applied  to  gases,  are  therefore  used  simply  because  it 
is  more  usual  to  see  gases  burning  in  oxygen  or  air  than  it  is  to 
see  oxygen  or  air  burning  in  other  gases.  Of  course,  the  phenom- 
ena attendant  upon  union  with  oxygen  also  appear  with  other 
gases  which  (chlorine,  for  example)  act  like  oxygen. 


CAIIBON    TETRACHLORIDE.  285 


CHAPTER   XXXIX. 

THE    COMPOUNDS    OP    CARBON    WITH    CHLORINE,    WITH   CHLO- 
RINE  AND   OXYGEN,   WITH   OXYGEN,   AND    WITH    SULPHUR. 

Carbon  dioxide  ;  formula,  C02 ;  specific  gravity,  air  =  1,  is  1.529, 
H2  =  2,  is  44  ;  molecular  weight,  44 ;  one  c.c.  of  the  gas  at  0°  and 
.76  m.  weighs  .001986  gram.  Carbon  monoxide  ;  formula,  CO  ; 
specific  gravity,  air  =  1,  is  .96744,  H2  =  2,  is  27.86  ;  molecular 
weight,  28 ;  1  c.c.  of  the  gas  at  0°  and  .76  m.  weighs  .0012511 
gram. 

THE  only  chloride  of  carbon  which  need  be  mentioned  in  this 
work  is  the  tetrachloride,  C  C14 ;  some  more  complicated  chlorine 
derivatives  of  carbon  chains  are  known,  but  a  work  on  organic 
chemistry  must  be  consulted  in  regard  to  their  properties.  Carbon 
tetrachloride  is  derived  from  methane  by  replacing  all  of  the  hydro- 
gen atoms  with  chlorine ;  and  it  can  be  prepared,  as  we  have  seen, 
from  the  latter  substance  by  the  action  of  chlorine ;  it  is  not,  how- 
ever, practically  expedient  to  commence  with  methane  in  order  to 
produce  the  tetrachloride,  because  methin  chloride  (C  H  C13 ,  chloro- 
form) is  easily  procured  by  other  means ;  and  then,  beginning  with 
this  chlorinated  methane,  we  can,  by  passing  chlorine  into  the  boil- 
ing liquid,  finally  substitute  the  remaining  hydrogen  atom  :  — 


Carbon  tetrachloride  is  a  colorless  liquid  which  boils  at  76°.5, 
and  which,  unlike  most  of  the  chlorides  of  the  not-metals,  is  but 
slowly  decomposed  by  cold  water  ;  on  warming  with  an  excess  of 
water  it  is,  however,  readily  converted  into  carbonic  and  hydro- 
chloric acids  ;  but  carbonic  acid,  like  all  acids  whose  anhydrides  are 
gases,  at  once  breaks  down  into  its  anhydride  and  water,  so  that, 
although  we  may  consider  orthocarbonic  acid  to  be  the  first  result  of 
the  reaction,  carbon  dioxide  is  the  only  tangible  carbon  compound 
produced  :  *  — 

*  Organic  derivatives  of  orthocarbonic  acid;  namely  orthocarbonic  acid, 

rOC2H5 


in  which  all  of  the  hydrogen  atoms  are  replaced  by  ethyl,  C 


286  CARBON   MONOXIDE. 

r  ci  +  HOH          r  OH 

Cl       HOH 


Cl  +  HOH  OH 

2.  C  (OH  )4  =  CO  (OH  )2  +  Ho  0,  and 

3.  CO  (OH  )2  =  C02  +  H2  0. 

The  action  of  alkalies  differs  from  that  of  water,  for  when  carbon 
tetrachloride  is  treated  with  potassium  or  sodium  hydroxide,  the 
stable  potassium  or  sodium  carbonate  is  produced. 

Carbontetrabromide,  C  Br4  ,  and  tetra-iodide,  C  I4  ,  are  also  known  ; 
the  former  is  a  solid,  which  melts  at  92°,  and  boils  at  189°.  5  ;  the 
latter  is  a  solid,  which  breaks  down  into  carbon  and  iodine  when 
heated. 

The  two  important  oxides  of  carbon  are  carbon  'monoxide,  CO, 
and  carbon  dioxide,  C02  ;  only  the  latter  acts  like  the  anhydride  of 
an  acid.* 

Carbon  dioxide  was  formerly  supposed  to  be  the  only  oxide  of 
carbon  ;  for  carbon  monoxide,  even  as  recently  as  the  beginning  of 
this  century,  was  believed  to  be  identical  with  hydrogen,  or,  at  least, 
to  contain  hydrogen.  Woodhouse,  of  Philadelphia,  in  the  year  J  800, 
first  proved  that  the  combustible  gas  obtained  by  reducing  metallic 
oxides  with  charcoal  was  not  hydrogen,  and  demonstrated  that  it 
contained  carbon;  but  carbon  monoxide  was  not  recognized  as  a 
combustible  oxide  of  carbon  until  after  the  year  1802.  | 

Carbon  monoxide  is  produced  by  the  incomplete  combustion 
and  also  by  the  dry  distillation  of  bituminous  coal  and  of  organic 
matter  ;  for  this  reason  it  occurs  in  illuminating  gas.  Carbon  mon- 
oxide is  likewise  always  formed  when  reducible  metallic  oxides,  such 
as  those  of  iron  or  zinc,  are  heated  with  charcoal  :  — 

ZnO  +  C  =  Zn  +  CO. 


The  heating  of  metallic  oxides  with  charcoal  is  a  general  method 
of  preparing  metals  from  their  ores.  These  reactions  are,  how- 
ever, like  many  others,  reversible,  so  that  carbon  dioxide  is  re- 
duced to  carbon  monoxide  by  metals  such  as  iron  or  zinc,  when 

*  A  few  cases  in  which  carbon  monoxide  exhibits  acidic  properties  are 
known.  For  instance,  it  unites  with  solid  caustic  potash  when  heated,  form- 
ing potassium  formate  :  CO  +  KOH  =  CHO2K. 

t  In  1801  Cruikshank  demonstrated  that  carbon  monoxide  is  an  oxide  of 
carbon. 


CARBON   MONOXIDE.  287 

these  are  heated  to  a  high  temperature.  At  1300°  carbon  monoxide 
is  partially  decomposed  into  carbon  dioxide  and  carbon.  When 
steam  is  passed  over  red-hot  charcoal,  carbon  monoxide  and  hydro- 
gen are  produced :  — 

H20  +  C=CO  +  2H, 

and  the  mixture  of  combustible  gases  so  obtained,  after  being 
passed  through  volatile  hydro-carbons,  is  quite  extensively  used  as 
illuminating  gas.*  Carbon  dioxide  can  also  enter  into  a  similar 
reaction ;  for,  when  it  is  passed  through  a  layer  of  hot  coal,  charcoal, 
or  coke,  carbon  monoxide  results,  as  is  shown  by  the  following 
equation  :  — 70 

C02  +  C=2CO. 

It  is  for  this  reason  that  the  carbon  dioxide,  produced  by  the 
free  combustion  of  the  coal  just  above  the  grate  in  a  stove,  is 
changed  to  carbon  monoxide  by  passing  over  the  hot  coals  above. 
In  many  cases  carbon  monoxide  acts  as  a  reducing  agent ;  for 
instance,  the  gas  passed  over  red-hot  ferric  oxide  reduces  the  latter 
to  metallic  iron  :  — 

Fe2 03  +  3  CO  =  2  Fe  +3  C02 . 

In  order  to  prepare  carbon  monoxide  for  laboratory  use  advan- 
tage is  taken  of  the  decomposition  of  oxalic  acid  by  heat  or  by 
means  of  concentrated  sulphuric  acid.f  The  acid  breaks  down 
as  follows  :  — 

CO  OH      co    :OH 

=rwvTT     °r,C204H2  =  C02  +  CO  +  H20.t 
CO  OH 

The  sulphuric  acid,  which  is  added,  assists  the  operation  by  reason 
of  its  great  tendency  to  take  up  water.71  The  carbon  dioxide,  which 
is  formed  simultaneously  with  carbon  monoxide,  can  be  removed 
by  passing  the  mixture  of  gases  through  a  solution  of  potassium 
hydroxide,  by  which  means  potassium  carbonate  is  formed,  while 
the  carbon  monoxide  can  be  collected  over  water. 

Carbon  monoxide  is  a  colorless  gas  which,  when  it  is  pure,  has 
scarcely  any  odor.  Its  specific  gravity,  air  =  1 ,  is  .96744.  Carbon 

*  So-called  water  gas. 

t  This  reaction  is  common  to  a  number  of  other  dibasic  organic  acids, 
t  Oxalic  acid  can  be  derived  from  ethane  by  replacing  four  of  the  hydro- 
gen atoms  in  one  molecule  by  oxygen  and  the  remaining  two  by  hydroxyl. 


288  CARBON  MONOXIDE;   PROPERTIES. 

monoxide  is  one  of  the  gases  which  is  with  difficulty  condensed 
to  a  liquid,  it  does  not  become  fluid  at  —  136°  and  under  150  atmos- 
pheres pressure;  however,  the  application  of  a  still  greater  cold 
changes  it  to  a  colorless  liquid  which  boils  at  —  190°  under  760  m.  m. 
pressure,  and  which  becomes  solid  at  —  207°.  One  volume  of  water 
dissolves  about  .023  of  a  volume  of  carbon  monoxide  ;  the  gas  is,  how- 
ever, quite  soluble  in  a  hydrochloric  or  an  ammoniacal  solution  of 
cuprous  chloride. 

Carbon  monoxide  burns  readily  in  oxygen  or  in  air  ;  the  product 
of  the  combustion  is  carbon  dioxide  :  — 


The  pale  blue  flames  observed  above  an  anthracite  coal  fire  are 
caused  by  carbon  monoxide. 

The  gas  acts  as  a  poison  ;  it  replaces  the  oxygen  which  is  chemi- 
cally combined  in  the  blood  by  an  equal  volume  of  carbon  monox- 
ide, each  molecule  of  carbon  monoxide  must,  therefore,  take  the 
place  of  a  molecule  of  oxygen  in  oxy  haemoglobin  ;  the  oxidizing 
powers  of  the  blood  are  thereby  destroyed,  for,  as  carbon  monoxide 
forms  a  more  stable  compound  with  haemoglobin  than  does  oxygen, 
it  is  obvious  that,  once  the  carbon  monoxide-haemoglobin  is  formed, 
this  compound  cannot  ,be  broken  up  by  oxygen.  Blood  which  has 
been  saturated  with  carbon  monoxide  retains  its  red  color  on  ex- 
posure to  the  air  for  a  longer  time  than  that  which  has  been 
oxygenated.* 

Carbon  monoxide,  when  mixed  with  chlorine  and  placed  in  the 
sunlight,  unites  directly  with  that  element  to  form  carbonyl  chloride 
(phosgen):  — 

CO  +  2C1  =  COC12; 

this  reaction  being  exactly  like  the  similar  one  observed  in  the  case 
of  sulphur  dioxide  (see  page  138).  Carbonyl  chloride,  at  ordinary 
temperatures,  is  a  colorless  gas  with  a  most  peculiar,  penetrating 
odor.  By  means  of  snow  and  salt  it  can  be  condensed  to  a  liquid 
which  boils  at  8°.  Water  readily  decomposes  carbonyl  chloride, 
forming  hydrochloric  acid  and  carbon  dioxide.  The  reaction  can  be 
considered  as  taking  place  in  two  phases  :  — 

*  Carbon  monoxide  can  readily  be  detected  in  the  blood  by  means  of  the 
peculiarity  of  the  absorption  spectrum  of  blood  saturated  with  the  gas. 


CARBON   DIOXIDE;    OCCURRENCE/  289 

/ 

f  Cl  +  HOH  (  OH      / 

1       C-}0  =     QJO     -f/^  H  Cl. 

(  Cl  +  HOH  (  OH 

f  OH 

(OH 

The  substance  reacts  in  a  similar  way  with  ammonia,  forming  car- 
bonyl  diamide  (urea)  and  hydrochloric  acid :  *  — 

Cl       NH8  ( NH2 

0  +  =     C-3  0      +2HC1. 

Cl       NH3  ( NH2 

A  study  of  these  reactions  shows  us  that  carbonic  acid,  H2  C03,  can 
be  considered  as  carbonyl  chloride  in  which  two  chlorine  atoms  in 
each  molecule  have  been  replaced  by  hydroxyl  groups ;  the  acid  itself, 
however,  does  not  exist ;  it  is  only  known  as  its  anhydride  C02 . 
From  the  above  it  is  also  evident  that  urea  is  carbonic  acid  in  which 
all  hydroxyl  groups  have  been  replaced  by  NH2 . 

Carbon  dioxide  is  of  far  greater  importance  than  carbon  monox- 
ide. Its  occurrence  in  the  atmosphere  and  the  manner  and  sources 
of  its  production  were  discussed  on  pages  166  and  167. 

Phenomena  by  which  carbon  dioxide  are  produced  were  known 
in  the  earliest  times,  but  the  gas  itself  escaped  observation.  It  was 
not  until  the  close  of  the  sixteenth  century  that  a  peculiar  gas, 
which  we  now  know  to  be  carbon  dioxide,  was  observed  escaping 
from  some  mineral  waters ;  Van  Helmont  (1577-1644)  first  distin- 
guished this  gas  from  others  and  gave  to  it  the  name  of  gas  sylvestre; 
he  showed  that  this  substance  was  produced  by  the  action  of  acids 
on  alkalies  or  lime,  by  the  burning  of  coals,  and,  in  addition,  was 
also  formed  during  the  processes  of  fermentation.  Black,  in  1757, 
showed  the  difference  between  the  so-called  caustic  alkalies  (now 
known  as  hydroxides)  and  mild  alkalies  (carbonates),  and  found 
that  a  peculiar  kind  of  air  (carbon  dioxide),  which  he  called  fixed 
air,  was  expelled  from  the  latter  by  the  addition  of  acids.  Lavoisier 
first  explained  the  true  nature  of  carbon  dioxide,  and  gave  to  it  the 
name  of  acide  carbonique. 

Pure  carbon  dioxide  can  best  be  prepared  by  the  addition  of  an 
acid  to  some  carbonate :  — 

*  The  monovalent  group  NH2  is  called  the  amido  group ;  see  page  192. 


290  CARBON   DIOXIDE;    PROPERTIES. 

DHa,  C03  +  2  H  Cl  =  2  Na  Cl  +  H,  0  -f  C02, 
Ca  C03  +  2  H  Cl  =  Ca C12  +  Ho  0  +  C02, 
Na^  C03  +  H2  S04  =  Na2  S04  -f  H2  0  +  C02. 

Because  carbonic  acid  is  one  of  the  weak  acids,  and  because  it  so 
readily  breaks  down  into  water  and  gaseous  carbon  dioxide  (see 
page  285),  it  follows  that  almost  any  other  acid  will  liberate  carbon 
dioxide  from  the  carbonates,  so  that  the  general  formula :  — 

M2  C03  +  2  H  X  =  2  M  X  +  H2  0  +  C02, 

(where  M  represents  a  univalent  metal,  and  H  X  a  monobasic  acid) 
will  hold  good  with  very  few  exceptions.*  The  most  convenient 
method  of  preparing  carbon  dioxide  for  laboratory  use,  is  by  the 
action  of  dilute  hydrochloric  acid  on  marble  (carbonate  of  calcium).72 
Carbon  dioxide  is  a  colorless  gas,  which  neither  burns  nor  sup- 
ports combustion.  It  has  a  specific  gravity  of  1.529,  air  =  1,  or  44, 
H2  =  2  ;  its  molecular  weight  is  therefore  44,  and  its  formula  C02 . 
This  formula  is  further  sustained  by  the  fact  that  there  is  no  change 
in  the  volume  of  the  gas  when  carbon  burns  in  pure  oxygen,  so  that 
each  molecule  of  carbon  dioxide  must  contain  one  molecule,  or  two 
atoms,  of  oxygen.  Because  carbon  dioxide  has  such  a  high  specific 
gravity,  it  can  be  poured  downward  from  any  vessel  containing  it, 
and  it  is  for  this  reason  that  carbon  dioxide  collects  at  the  bottom 
of  wells  and  mines  into  which  the  gas  is  escaping.73  Cold  and  pres- 
sure combined  condense  carbon  dioxide  to  a  liquid  which  boils  at 
— 78°. 2  ;f  the  vapor  tension  of  fluid  carbon  dioxide  is  36  atmospheres 
at  0°,  and  73  atmospheres  at  30°;  the  critical  point  is  30°. 9 ;  above  this 
temperature  no  pressure  can  convert  the  gas  into  a  fluid.  When 
carbon  dioxide  rapidly  evaporates  in  a  vacuum,  the  temperature 
sinks  to  —  97° ;  the  liquid,  when  allowed  to  escape  from  a  small 
opening,  condenses  to  a  white,  snow-like  mass,  the  temperature  of 
which,  at  atmospheric  pressure,  is  — 78°.  Liquid  carbon  dioxide  is 
colorless,  and  has  a  specific  gravity  of  .995  at  10°.  Liquid  carbon 
dioxide  is  extensively  used  in  commercial  operations ;  for  instance, 

*  Some  few  acids,  like  hydrocyanic  acid,  are  unable  to  decompose  carbon- 
ates, while  some  few  carbonates  which  occur  as  minerals  (for  instance,  dolo- 
mite) are  not  readily  attacked  by  dilute  acids.  Quite  a  number  of  organic 
substances  which  act  like  acids  are  unable  to  decompose  carbonates. 

t  Doubtful.  The  temperature  of  melting  carbon  dioxide,  mixed  with 
ether,  is  —80°  (Dewar  and  Fleming;  Philosph.  Mag.;  34,  329). 


CARBONATES,    PRIMARY.  291 

in  the  manufacture  of  soda  water,  in  fire  extinguishers,  and  in  oper- 
ations where  the  great  pressure  exerted  by  it  can  be  used.  It  is 
transported  in  thick-walled  steel  tubes. 

Carbon  dioxide  is  the  anhydride  of  carbonic  acid,  but  the  latter 
substance  is  extremely  unstable.  It  is  probably  formed  as  a  white 
mass  when  the  pressure  is  suddenly  removed  from  carbon  dioxide 
which,  in  the  presence  of  water,  has  been  nearly  condensed  to  a 
liquid  at  0°.  There  is  no  probability  that  water  which  is  saturated 
with  carbon  dioxide  at  ordinary  temperatures  contains  the  acid 
H2  C03  as  such,  for  the  solution  behaves  physically  like  an  ordinary 
gas  solution,  and  not  like  that  of  an  acid ;  on  the  other  hand,  the 
carbonates  of  the  pronounced  metals  are  extremely  stable  sub- 
stances. With  a  diminution  of  the  metallic  nature  of  the  salt- 
forming  element,  the  -carbonates  become  less  stable,  and  very  weak 
bases,  like  the  oxide  of  aluminium  or  ferric  oxide,  cannot  react  with 
carbonic  acid  at  all. 

The  carbonates  are  all  derived  from  a  dibasic  acid  H2C03.'  The 
secondary '  carbonates,  M2C03,  are,  as  a  rule,  insoluble  in  water ; 
only  those  of  the  alkali  metals  and  of  ammonium  *  dissolve ;  the 
other  carbonates  can  therefore,  be  obtained  from  these  .by  precipita- 
tion with  the  soluble  salt  of  some  other  metal,  for  example :  — 

Na*  C03  +  Ba  C12  =  Ba  C03  +  2  Na  Cl. 

Soluble.  Soluble.       Insoluble.     Soluble. 

The  carbonates  of  the  alkali  metals  can  be  fused  without  change ; 
all  other  carbonates  are  more  or  less  readily  decomposed  into  car- 
bon dioxide  and  the  metallic  oxide  by  heating,  thus :  — 

CaC03=  CaO  +  C02, 

and  this  decomposition  takes  place  the  more  readily,  the  less  basic 
the  metallic  oxide  is,  so  that  many  carbonates  are  even  decomposed 
on  boiling  with  water.  This  increasing  stability  of  the  carbonates 
with  the  increase  in  the  metallic  character  of  the  salt-forming  ele- 
ment, is  exactly  parallel  to  the  same  gradation  observed  in  the 
chlorides  of  the  elements  of  the  phosphorus  family  (see  page  181). 

The  primary  carbonates,  with  the  exception  of  those  of  the 
alkalies,  exist  only  in  aqueous  solution ;  they  can  be  obtained,  where 
their  existence  is  possible,  by  treating  a  solution  of  a  secondary 
carbonate,  or  even  a  finely  divided  insoluble  secondary  carbonate 

*  Lithium  carbonate  is  soluble  with  difficulty. 


292  CARBONATES;  SECONDARY. 

suspended  in  water,  with  carbon  dioxide  ;  they  are  unstable  and  are 
readily  broken  down'  by  heat :  — 

2  Na  HC03  =  Na2  C03  -f  H2  0  +  C02 ; 

where  they  exist  they  are  soluble.  The  solution  of  calcium  car- 
bonate in  temporary  hard  water  is  caused  by  the  formation  of  the 
primary  calcium  carbonate  by  means  of  the  carbon  dioxide  con- 
tained in  the  air  or  added  to  the  water  by  decaying  substances  :  — 

CaC03  +  H2C03  =  Ca  (HC03)2. 

This  soluble  primary  carbonate  is  decomposed  when  the  water 
evaporates  *  or  when  it  is  heated.  The  temporary  hard  waters  for 
this  reason  deposit  their  calcium  carbonate  as  a  white  coating  on 
the  walls  of  the  kettle  in  which  they  are  boiled. 

Secondary  carbonates,  when  soluble,  have  a  strongly  alkaline 
reaction,  the  primary  ones  are  neutral. f 

The  following  table  gives  a  few  of  the  most  important  naturally 
occurring  carbonates :  — 

Calcium  carbonate   )  Massive  varieties;  chalk,  limestone,  marble. 

(CaCO)  >  Crystallized  varieties;   calcite  (Iceland  spar),  arra- 

gonite. 

Calcium  and  magnesium  carbonate  (Ca,  Mg),  CO3,  dolomite. 
Ferrous  carbonate;  Fe  CO3  ,  siderite. 
Barium  carbonate;  Ba  CO3,  witherite. 
Strontium  carbonate ;  Sr  CO3  ,  strontianite. 

The  above  carbonates  are  frequently  found  as  isomorphous  mix- 
tures. The  carbonates  of  lead,  zinc,  and  manganese  are  also  found, 
as  well  as  basic  carbonates  of  copper,  bismuth,  and  zinc.  Carbonates 
of  sodium  with  more  or  less  water  of  crystallization  occur  as  soda 
(  Na2  C03  +  10  H2  0)  and  trona  (  Na4  H2  [C03  ]3  3  H2  0) .  t 

The  compound  of  carbon  and  sulphur  which  corresponds  to  C02 
is  CS2,  carbon  disulphide.  This  liquid  can  be  formed  by  heating 
carbon  in  sulphur  vapor,  so  that  the  method  of  its  production  cor- 
responds to  that  of  carbon  dioxide.  Carbon  disulphide  is  a  colorless, 
mobile  liquid,  which,  when  pure,  has  a  pleasant  ethereal  odor.  Its 
specific  gravity  is  1.29  ;  the  specific  gravity  of  its  vapor  is  2.626 
(air  =  1)  ;  it  boils  at  48°,  and  is  very  little  soluble  in  water.  When 

*  Formation  of  stalactites. 

t  They  have  no  effect  on  litmus  or  turmeric  paper,  but  do  have  an  alkaline 
reaction  toward  rosolic  acid. 

t  (2  Na  HC03  +  Na2  CO3  +  3 II,  O. ) 


THIOCARBONATES.  293 

gently  warmed  with  the  alkaline  sulphides,  carbon  disulphide  is 
dissolved  while  the  salts  of  sulpho-acids  are  formed  :  — 


These  salts  of  trithiocarbonic  acid  (sulpho-dithio  carbonic  acid)  * 
correspond  to  those  of  carbonic  acid,  the  oxygen  atoms  in  the  latter 
having  been  replaced  by  sulphur.  When  acids  are  added  to  its 
salts,  sulpho-dithio  carbonic  acid,  H2CS3,  separates  as  an  oil:  — 

K2CS3  +  2HC1==H2CS3  +  2KC1, 

but  the  thio  acid  so  produced,  although  it  is  not  as  unstable  as 
the  corresponding  oxy-acid,  nevertheless,  gradually  breaks  down  as 

follows  : 

H2CS3=H2S  +  CS2, 

just  as  carbonic  acid  decomposes  into  water  and  carbon  dioxide. 
These  reactions  of  carbon  disulphide  remind  us  forcibly  of  the  simi- 
lar ones  encountered  with  the  sulphides  of  arsenic  and  antimony. 
(See  pages  245,  255.)  A  compound  of  carbon,  oxygen,  and  sulphur, 
having  the  formula  COS  (lying  between  the  dioxide  and  sul- 
phide of  carbon),  as  well  as  acids  derived  from  it,  are  also  known  ; 
and  several  sulphides  of  carbon,  differing  from  carbon  di-sulphide 
(for  example,  C2  S3  )  have  also  been  described. 

*  The  name  trithiocarbonic  acid  is  derived  from  Qftov,  sulphur,  and  the 
name  thio-acids  is  frequently  employed  for  sulpho-acids  and  thio-compounds 
for  sulpho-eompounds.  An  endeavor  is  made  to  establish  the  following  dis- 
tinction: where  sulphur  is  attached  to  carbon  only,  it  is  called  sulpho,  where 
it  is  attached  to  carbon  on  the  one  hand  and  a  metal,  or  a  group  of  elements 

<—  SH    ^  CL 

acting  like  a  metal,  on  thej>ther,  it  is  called  thio;  thus:  C  j  O   /    is(   dttkiit 


(  —  SH 

carbonic  acid,  and  C  ]  S_       is  sulphothioc&vltomc  acid.      This   nomenclature 
I—  OH  *    B' 


is,  however,  logically  carried  out  only        h  tne  compoundsjjkf  carbon. 

* 


sjjkf  c 


294  CYANOGEN. 


CHAPTER   XL. 

COMPOUNDS  OF  CARBON  WITH  NITROGEN,  WITH  NITROGEN 

AND  HYDROGEN,  AND  WITH  NITROGEN,  OXYGEN, 

AND  HYDROGEN. 

ONLY  a  very  few  of  the  more  important  of  these  compounds  can 
be  strictly  considered  as  belonging  to  the  realm  of  inorganic  chem- 
istry, and  these  only  will  be  briefly  considered  in  this  work.  The 
most  prominent  of  the  substances  to  be  discussed  are  derivatives  of 
the  monovalent  group  of  elements,  cyanogen,*  CN.  This  group  can 
be  attached  to  other  elements  or  groups  of  elements  in  two  ways, 
either  through  the  element  nitrogen,  by  which  means  substances 
having  the  general  structural  formula  MNC  are  formed ;  or  through 
carbon,  the  general  structural  formula  of  the  latter  class  of  com- 
pounds being  MCN ;  in  the  first  case  isocyanides  are  formed,  in  the 
second  true  cyanides  (also  called  nitriles).  Representatives  of  both 
classes  of  compounds  are  known,  f 

All  nitrogen-bearing  organic  compounds,  or,  indeed,  the  nitro- 
genous coals  derived  from  these,  yield  the  cyanide  of  sodium  or  of 
potassium  when  heated  with  one  or  the  other  of  those  metals ;  the 
cyanide  of  potassium  is  also  formed  wh.en^coal  which  contains  nitro- 
gen compounds  is  heated  with  common  pbtaSff ;  by  this  process  the 
greater  quantity  of  the  cyanide  of  potassium  which  finds  commer- 
cial %application  is  prepared.  Free  cyanogen  can  be  formed  by 
heating  the  cyanide  of  mercury  to  a  dark  red  heat :  — - 


The  group  CN  is  as  incapable  of  individual  existence  as  is 
methyl,  CH3;  two  of  the  radicles  therefore  unite  to  form  dicya- 
nogen,  CN  —  CN,  just  as  methyl  unites  to  form  dimethyl.  | 

*  From  xi'avo'>i  blue. 

t  The  pupil  must  refer  to  some  larger  work  on  organic  chemistry  for  the 
reactions  which  characterize  the  isocyanides. 

J  Cyanogen  can  be  considered  as  methyl  in  which  three  atoms  of  hydro- 
gen are  replaced  by  one  atom  of  trivalent  nitrogen;  viz., 


HYDROCYANIC    ACID.  295 

Cyanogen  is  a  colorless  gas,  extremely  poisonous,  with  an  irri- 
tating odor.  Its  specific  gravity  is  1.804,  air  =  1,  or  51.9,  H  =  2, 
while  the  molecular  weight  of  (CN  )2  is  52.  The  gas  is  quite  solu- 
ble in  water,  and  is  combustible,  burning  with  a  characteristic 
purple  flame.  It  liquefies  at  —  20°.7  and  becomes  solid  at  —  34°.  4. 
Aqueous  solutions  of  cyanogen  gradually  decompose. 

Chemically,  cyanogen  greatly  resembles  the  halogens  ;  the  mono- 
valent  group  GN",  when  united  with  hydrogen,  forms  hydrocyanic 
acid,  just  as  chlorine,  similarly  united,  forms  hydrochloric  acid; 
furthermore,  hydrogen,  sodium,  or  potassium  will  unite  directly 
with  cyanogen  to  form  hydrocyanic  acid,  sodium  cyanide,  or  potas- 
sium cyanide,  just  as  the  same  elements  will  unite  with  chlorine  to 
form  hydrochloric  acid,  sodium,  or  potassium  chloride.  The  group 
CNj  cyanogen,  is  therefore  a  compound  radicle  acting  like  a  not- 
metal,  and  is  chemically  the  opposite  of  the  metal-like  radicle 
ammonium. 

Hydrocyanic  acid  *  is  prepared,  according  to  the  general  method, 
by  adding  an  acid  to  a  cyanide  :  — 

2  KCN  +  H2  S04  =  2  HCN  +  K2  S04  .f 

Or,  better,  by  distilling  commercial  potassium  ferrocyanide  with 
diluted  sulphuric  acid  (concentrated  acid  cannot  be  used  for  this 
purpose,  as  it  generates  carbon  monoxide).  When  pure,  hydrocy- 
anic acid  is  a  colorless,  mobile  liquid,  which  boils  at  26°  and  melts 
at  —  14°  ;  it  is  without  acid  reaction  toward  litmus,  and  has  a 
peculiar  odor,  somewhat  resembling  that  of  bitter  almonds.  It  is 
intensely  poisonous  even  when  inhaled  in  small  quantities  ;  one 
drop  placed  on  the  tongue  of  a  dog  will  cause  instant  death  ;  the 
poison  can  also  act  through  contact  with  abrasions  of  the  skin. 
The  acid  mixes  with  water  in*all  proportions,  and  solutions  of  vary- 
ing strength  form  the  commercial  prussic  acid  ;  the  aqueous  solution 
decomposes  on  standing.  The  acid  was  first  prepared  in  a  pure 
state  by  Gay  Lussac,  in  1811. 


Compare  the  preparation  of  cyanogen  with  the  preparation  of  oxygen  by  heat- 
ing mercuric  oxide. 

*  Also  called  prussic  acid  because  it  is  the  source  of  prussian  blue. 

t  The  usual  method  is  to  decompose  potassium  ferrocyanide  (see  latter) 
with  sulphuric  acid. 


296  POTASSIUM   FERROCYANIDE. 

Two  structural  formulae  are  possible  for  hydrocyanic  acid; 
namely,  C  =  N  —  H  and  H  —  C  =  N  ;  in  the  first  one  of  these  the 
hydrogen  atom  is  attached  to  nitrogen,  in  the  second  to  carbon  ; 
organic  derivatives  of  both  forms  are  known  ;  the  cyanides  of  the 
metals  are  most  probably  derivatives  of  the  first  form,  CNH.* 

Hydrocyanic  acid  is  a  very  weak  acid,  the  cyanides  of  the 
alkali  metals  and  of  barium,  calcium,  or  strontium  are  even  decom- 
posed by  moist  carbon  dioxide,  so  that  these  substances,  when  in 
contact  with  '  the  air,  emit  the  odor  of  hydrocyanic  acid.  The 
cyanides  of  most  of  the  heavy  metals  are  insoluble,  an  exception 
to  this  rule  is  the  cyanide  of  mercury  ;  some  of  these  cyanides  are 
not  readily  decomposed  by  acids  ;  t  on  the  other  hand,  the  alkaline 
cyanides,  which  are  so  readily  decomposed  by  weak  acids,  are  ex- 
tremely stable  when  heated  ;  they  can  even  be  fused  without  under- 
going a  chemical  change.  Almost  all  cyanides  which  are  soluble  in 
water  are  converted  into  so-called  double  cyanides  when  treated 
with  the  cyanides  of  the  alkalies.  Many  of  these  double  cyanides 
are  stable,  crystalline  bodies  which  have  the  nature  of  chemical 
compounds  Q  for  example,  ferrous  cyanide  when  brought  in  contact 
with  potassium  cyanide  forms  a  double  cyanide  of  the  formula 


Fe(CN)a  +      4KCN  =      K4(CN)6Fe. 

Ferrous  cyanide        +        Potassium  cyanide        =        Potassium  ferrocyanide. 

Potassium  ferrocyanide  bears  no  resemblance  either  to  potassium 
cyanide  or  to  ferrous  cyanide  ;  it  is,  indeed,  the  salt  of  a  tolerably 
strong  crystalline  acid,  hydroferrocyanic  acid,  H4(CN)6Fe.  This 
acid  can  be  prepared  by  the  addition  of  strong  hydrochloric  acid  to 
potassium  ferrocyanide;  the  ferric  salt  of  this  acid,  ferric  ferrous 
cyanide,  can  be  obtained  from  potassium  ferrocyanide  by  adding  a 
ferric  salt  :  — 

3  K4  (CN  )6  Fe  +  4  Fe  C13  =  Fe4  [(CN  )6  Fe]3  +  12  K  Cl. 
This  substance  is  the  insoluble  blue  dye  known  as  prussian  blue. 

Ferric  cyanide  is  also  able  to  form  a  double  salt  with  potassium 
cyanide  :  — 

*  The  experiments  which  seem  to  indicate  that  cyanides  are  derived  from 
both  forms  of  hydrocyanic  acid  can  easily  be  explained  on  the  assumption 
that  they  are  all  derived  from  the  first  form.  Compare  Nef  ;  Liebig's  Anna- 
len;  270,  329. 

t  Cyanides  of  mercury,  silver,  and  gold. 


POTASSIUM  FERRICYANIDE  ;   CYANIC   ACID.  297 

Fe  (CN  )3  +  3  KCN  =  K3  (CN  )6  Fe, 

and  this  substance,  on  addition  of  a  ferrous  salt,  also  forms  an  in- 
soluble blue  compound,  Turnbull's  blue.*  The  cyanides  of  other 
metals  which  are  chemically  closely  allied  to  iron  can  form  similar 
double  cyanides  ;  a  larger  work  must  be  consulted  in  regard  to  the 
properties  of  those  compounds. 

When  the  cyanide  of  potassium  is  oxidized,t  it  changes  to  the 
cyanate  of  potassium,  CNOK.  This  oxidation  is  easily  explained  if 
we  consider  hydrocyanic  acid  to  be  analogous  to  carbon  monoxide  :  — 

C=N  —  H,  C=0, 

Hydrocyanic  acid,  Carbon  monoxide  ; 

for  then  it  would  naturally  follow  that  the  former  would  be  oxidized 
as  readily  as  is  the  latter  :  — 


C=N—  H  +  0  =  C  |  ^T~H     and  C=0  +  0  =  C  j 


It  seems  highly  probable,  however,  that  the  salts  of  this  cyanic 
acid  assume  another  form,  namely,  one  in  which  the  metallic  atom 
is  attached  to  oxygen  :  —  t 


j  =  N  —  H   ,   0  —  H  J^N 

1  =  0  'K       ,  =    °|-0. 


K 

Such  a  difference  between  the  structure  of  the  free  acid  and  of 
the  salts  derived  from  it  seems  not  unlikely,  if  the  great  tendency 
which  is  displayed  by  potassium  or  sodium  to  unite  with  oxygen  is 
remembered.  The  acid  having  the  formula  CONH  is  termed  isocy- 
anic  acid,  and  the  one  with  the  structure  CNOH,  normal  cyanic  acid. 
Ordinary  cyanic  acid  is  at  present  known  only  in  one  form  ;  it  is  a 
very  mobile,  volatile  liquid,  which  is  stable  only  below  0°,  its  odor 
resembles  that  of  sulphur  dioxide.  A  polymeric  form  of  cyanic 
acid  (CNOH  )3 ,  cyanuric  acid,  is  also  known,  it  is  a  solid  substance.  § 

*  Probably  identical  with  prussian  blue. 

t  This  oxidation  even  takes  place  upon  exposing  potassium  cyanide  to  the 
air. 

J  Organic  derivatives  obtained  from  the  cyanates  undoubtedly  are  derived 
from  cyanic  acid  of  the  first  formula,  CONH ;  but  this  fact  does  not  seem  to 
prove  that  the  metal  in  the  cyanates  is  attached  to  nitrogen. 

§  It  cannot,  as  yet,  be  absolutely  decided  which  of  the  two  forms,  CNOH 
or  CONH  is  represented  by  ordinary  cyanic  acid.  Organic  derivatives  of 
both  acids  (those  in  which  hydrogen  is  replaced  by  ethyl,  etc.)  are  known, 


298  UREA;   CARBAMIC   ACID. 

Ammonium  cyanate,  on  standing,  or  more  rapidly  on  heating, 
changes  into  urea,  a  substance  of  the  greatest  physiological 
importance :  -  CNONH4  =  CO  (  NH2 ), 

Urea  can  be  considered  as  carbonic  acid  in  which  both  hydroxyls 
have  been  replaced  by  the  amido  group,  NH2 . 

(OH  (  —  NH2 


(OH  (—  NH2 

Carbonic  acid.  Urea. 

A  compound  derived  from  carbonic  acid,  but  in  which  but  one 
hydroxyl  group  is  replaced  by  the  amido  group  (NH2)  is  also 
known ;  this  substance  is  termed  carbamic  acid,  C02  HNH2 :  — 

(0  ro 

C  <  OH  carbonic  acid  and  C  -j  OH  carbamic  acid. 
(  OH  (  NH2 

The  ammonium  salt  of  this  acid, 

^-\ 

is  found  in  commercial  carbonate  of  ammonium ;  on  heat- 
ing to  130°  it  changes  into  urea  :  — 

0  CO 


NH2 

Another  interesting  method  by  which  urea  can  be  formed  is  by 
the  action  of  ammonia  on  carbonyl  chloride  :  — 

( Cl       NH3  ( NH2 

(VO+  =   CJ  0       +  2HC1, 

( Cl       NH8  ( NH2 

for  this  change  is  analogous  to  the  reaction  which  takes  place  when 
water  acts  on  the  same  substance  :  — 

(  Cl  +  HOH  (  OH 

C^O  =(OO+  2  HC1  (see  page  289). 

(  Cl  +  HOH  (  OH 

The  oxygen  atom  in  a  formula  weight  of  urea  can  be  replaced  by 
the  divalent  group  =  NH,  and  the  resulting  compound  (guanidine), 
which  has  the  formula  :  — 

although  those  of  normal  cyanic  acid  (CNOH  )  have  never  been  obtained  in  a 

I  =  O 
pure  state.    It  seems  not  unlikely  that  the  free  acid  has  the  structure  C     ~ 

while  the  salts  are  derived  from  C 


GUANIDINE  ;    CYANAMIDE. 


299 


(HH,, 

differs  but  little  from  urea  in  its  character. 

The  foregoing  compounds  are  extremely  interesting  from  a  theo- 
retical standpoint,  because  they  illustrate  the  close  chemical  resem- 
blance between  the  amido  and  hydroxyl  groups ;  for  the  former  can 
take  part  in  the  formation  of  compounds  which  are  similar  in  char- 
acter to  the  more  familiar  ones  derived  from  the  latter,  and  ammonia 
can,  therefore,  play  much  the  same  role  as  water  in  the  reactions 
entered  into  by  the  carbon  compounds  which  have  just  been  dis- 
cussed. This  similarity  between  nitrogen  and  oxygen  compounds 
serves  forcibly  to  illustrate  the  fact  that  no  element  is  isolated  in 
properties  and  chemical  character,  and  it  should  be  the  endeavor  of 
every  student  of  chemistry  to  detect  and  understand  the  resem- 
blances which  are  in  reality  present. 

The  connection  between  carbonic  acid  and  the  above  compounds 
is  shown  by  the  following  table  :  — 


2, 


«,   (  OH 

b.   (OH        c. 

(OH 

d.  (  OH          c.    r  NH2 

1        (0  OH 

L        "  1  OH 

r  J  OH         n 

Ut  j  OH 

j  OH 
J  NH 

C'  J  NHo        C'  j  NH2 

I  OH 

^NH2 

CKH 

2              I  NH2              I  NH2 

less  water  give, 

less  ammonia  give, 

'(° 

(0 

(0                           (  NH 

C  <  OH 

C  ]  OH 

C  ]  NH2                   C  ]  NH2 

(OH 

(  ^H2 

(  J^H2                      (  NH2 

carbonic  acid 

,               carbamic  acid, 

urea,                          guanidine, 

—  H2O 

-H20              —  N 

H3 

-H20         -NH3         -NI 

o          C(0 

cio      ^I^H  ^|o      "iNHo      ^INH 

carbon  dioxide,   cyanic  acid,  carbon  dioxide,  cyanamide,  cyanic  acid,    cyanamide. 

The  hypothetical  compounds  &,  c,  cZ,  e,  are  supposed  to  be  derived  from  the 
hypothetical  normal  carbonic  acid  (a)  by  replacing  hydroxyl  groups  by  amido 
groups  (OH  by  NH2)  .  In  this  table  the  hydroxyl  group  is  considered  as  anal- 
ogous to  the  amido  group  and  of  equal  valence,  and  the  group  NH  (imid 
group)  as  analogous  to  a  divalent  oxygen  atom.  Of  the  compounds  under  (2), 
carbonic  acid  and  carbamic  acid  are  known  only  in  their  derivatives. 


*  Cyanamide  is  a  solid,  melting  at  40°  and  differing  from  urea  by  one 
molecule  of  water. 


300  SILICON;    PREPARATION. 


CHAPTER   XLI. 

SILICON,  THE   COMPOUNDS  OF   SILICON  WITH  HYDROGEN, 

AND   WITH   THE   HALOGENS,  THE  OXIDE,  AND 

ACIDS   OF   SILICON. 

Silicon;  symbol,  Si  ;  atomic  weight,  28.4  ;  specific  gravity  of  solid 
(graphitoidal  ),  2.49. 

SILICON  never  occurs  in  nature  as  the  uncombined  element,  but 
silicon  compounds  are  among  the  most  important  and  widely  dis- 
tributed constituents  of  the  crust  of  the  earth.  The  primitive  crys- 
talline rocks  are  in  greater  part  either  silicon  dioxide  or  else  salts 
of  the  various  silicic  acids.  It  has  been  estimated  that  27.2  per 
cent  of  the  globe  (excluding  the  atmosphere)  consists  of  silicon. 
Despite  the  abundance  of  silicon  compounds,  the  element  itself 
was  not  discovered  until  1823,  in  which  year  it  was  isolated  by 
Berzelius. 

Silicon  is  best  prepared  by  the  reduction  of  some  of  the  halo- 
gen compounds  of  the  element  by  means  of  sodium  or  potassium  ; 
for  instance,  by  passing  the  vapors  of  the  tetrachloride  of  silicon 
over  heated  sodium  :  - 


SiCl4  +  4Na=Si  + 

The  element  so  prepared  is  an  amorphous  brown  powder,  which 
does  not  conduct  electricity  ;  it  is  readily  ignited  in  the  air,  burning 
to  form  silicon  dioxide.  Silicon  is  produced  when  potassium  or 
sodium  fluosilicates  *  are  fused  with  aluminium  ;  the  aluminium, 
uniting  with  the  fluorine,  in  part  forms  aluminium  fluoride,  while 
the  unchanged,  molten  metal  dissolves  the  silicon  which  is  liberated  ; 
when  the  mass  is  cooled  a  portion  of  the  silicon  separates  in  needle- 
shaped  crystals.f  Silicon  crystallizes  in  grayish  black,  regular 
octahedra  which  have  a  metallic  lustre.  The  silicon  which  is 

*  K2  Si  F6  or  Na2  Si  F6  ,  salts  of  fluosilicic  acid,  KU  Si  F6  . 

t  Melted  zinc  or  iron  can  also  dissolve  silicon  ;  when  they  cool,  the  silicon 
separates  in  crystals  ;  as  silicon  is  formed  from  silicon  dioxide,  carbon,  and  iron 
at  a  high  heat,  it  follows  that  ordinary  pig  iron  must  contain  silicon. 


SILICON  ;    HYDRIDE,   TETKACHLORIDE.  301 

formed  from  molten  iron  or  zinc  may  also  have  the  appearance 
of  graphite,  although  the  crystalline  form  is,  apparently,  the  same 
as  that  of  the  needle-like  crystals ;  graphitoidal  silicon  has  a  specific 
gravity  of  2.49 ;  crystallized  silicon  conducts  electricity  readily  and 
cannot  be  ignited  in  the  air.  When  heated  to  a  high  white  heat, 
silicon  can  be  fused  and  even  cast  into  sticks ;  when  cooled  it  solidi- 
fies to  form  a  mass  which  somewhat  resembles  a  piece  of  pure 
crystalline  graphite.  The  amorphous  and  crystalline  varieties  of 
silicon  remind  us  of  the  similar  forms  displayed  by  carbon. 

Only  one  compound  of  silicon  and  hydrogen,  silicon  hydride, 
Si  H4 ,  is  known.  This  substance  bears  a  great  resemblance  to  the 
corresponding  compound  in  the  nitrogen  family,  namely,  to  phos- 
phine,  for  it  takes  fire  spontaneously  when  brought  into  the  air, 
and  it  is  formed  by  the  action  of  an  acid  on  magnesium  silicide ; 
in  principle  this  is  the  same  method  as  that  employed  for  the  prepa- 
ration of  phosphine  from  calcium  phosphide.  The  preparation  of 
pure  silicon  hydride  is  a  difficult  process.  Silicon  hydride  is  a 
colorless  gas  which,  when  pure,  does  not  take  fire  spontaneously, 
but  which  has  such  a  low  kindling  temperature  that  it  can  readily 
be  ignited  by  a  warm  glass  rod ;  the  gas  is  liquefied  at  —  11°  by  a 
pressure  of  50  atmospheres.  Silicon  hydride  is,  of  course,  readily 
decomposed  by  chlorine  or  bromine,  and  it  resembles  hydrogen 
sulphide  and  phosphine  by  producing  precipitates  with  quite  a 
number  of  metallic  salts. 

The  compounds  of  silicon  with  the  halogens  possess  general  for- 
mulae similar  to  the  corresponding  compounds  of  carbon ;  silico- 
chloroform,  Si  H  C13,  corresponding  to  ordinary  chloroform,  CHC13, 
and  silicon  tetrachloride,  SiCl4,  corresponding  to  CC14,  serve  to 
illustrate  this  resemblance.  Silicon  tetrachloride  is  a  colorless, 
mobile  liquid  which  boils  at  58°,  and  which  is  energetically  de- 
composed by  water.  It  has  a  most  penetrating  and  irritating  odor 
which  resembles  that  of  phosphorus  pentachloride.  The  halogen 
compounds  of  silicon  are  all  unstable  bodies  which  are  decomposed  by 
water  to  form  silicic  acid  and  the  corresponding  halhydric  acid  :  — 

Si  C14  +  3  H20  =  Si  03  H2  +  4  H  Cl . 

The  most  important  halogen  compound  of  silicon  is  undoubtedly 
the  tetrafluoride,  Si  F4 .  This  substance  can  readily  be  produced 
by  the  action  of  hydrofluoric  acid  on  silicon  dioxide  :  — 


302  SILICON    AND    TETKAFLUOUIDE. 

Si  02  +  4  HF  =  Si  F4  +  2  H2  0 . 

In  this  reaction  the  silicon  dioxide  is  a  base,  for  it  yields  a  salt  and 
water  when  brought  in  contact  with  an  acid.  Silicon  tetrafluoride 
is  a  colorless  gas,  which  fumes  strongly  in  the  air,  because  when  in 
contact  with  water  vapor  it  decomposes  and  forms  silicic  acid ;  it 
has  a  most  penetrating  and  irritating  odor,  arid  it  reddens  litmus 
paper,  even  when  dry.  At  a  temperature  of  —  105°  and  at  9  atmos- 
pheres pressure  silicon  tetrafluoride  is  converted  into  a  colorless 
liquid,  which  becomes  solid  at  —  140°.  The  specific  gravity  of  the 
vapor,  air  =1 ,  is  3.6 ,  a  number  which  would  agree  with  a  calcu- 
lated molecular  weight  of  103.6,  so  that  according  to  this  the  mole- 
cule of  silicon  tetrafluoride  is  SiF4.  As  silicon  tetrafluoride  is 
readily  produced  by  the  action  of  hydrofluoric  acid  on  silicon 
dioxide,  it  follows  that  hydrofluoric  acid  will  attack  glass,  for  that 
substance  contains  a  large  proportion  of  silicon  dioxide.74  Water 
instantly  decomposes  silicon  tetrafluoride,  forming  silicic  acid  and 
fluosilicic  acid  :  — 

2  SiF4  +  3  H20  =  H2  Si  03  +  Si  F6H2  +  2  HF. 

Fluosilicic  acid  is  especially  interesting  because  it  shows  the  chem- 
istry of  fluorine  to  us  in  an  entirely  new  light ;  for,  if  we  compare 
the  following  two  formulae  :  — 

(0  (F2 

Si-}  OH         and         Si^F2H 
(OH  (F2H 

Silicic  acid          and  Fluosilicic  acid, 

we  see  that  fluosilicic  acid  is  constructed  similarly  to  silicic  acid, 
but  with  this  difference ;  for  each  oxygen  atom  in  silicic  acid  we 
have  two  fluorine  atoms  in  fluosilicic  acid.  Two  fluorine  atoms  in 
some  chemical  compounds  are  therefore  able  to  take  the  place  of 
one  oxygen  atom  without  materially  altering  the  chemical  nature 
of  those  compounds.  Substances  which  are  constructed  in  a  manner 
similar  to  fluosilicic  acid  are  not  infrequent ;  but,  as  their  resem- 
blance to  oxygen  compounds  is  not  generally  so  marked  as  in  the 
case  under  discussion,  their  true  nature  is  often  misunderstood  and 
concealed  under  the  names  of  double  salts  (see  aluminium).*  The 

*  Hydro-ferrocyanic  acid,  H4Fe  (CN)6,  may  be  cited  as  an  instance 
where  a  complex  compound  containing  hydrogen  (and  in  which  cyanogen  has 
taken  the  place  of  oxygen)  can  act  as  an  acid. 


FLUOSILICIC   ACID.  303 

wider  the  range  of  our  acquaintance  with  chemical  compounds  be- 
comes, the  more  do  we  see  that  the  most  various  substances,  which 
may  or  may  not  contain  oxygen,  can  act  as  acids,  provided  only 
they  contain  hydrogen  attached  to  a  not-metallic  element  or  group 
of  elements,  and  which  hydrogen  can  be  replaced  by  metals  to 
form  salts  (see  page  75).  The  reason  for  the  salt  formation  is 
found  in  the  simple  fact  that  the  salts  produced  by  the  replace- 
ment of  the  acid  hydrogen  possess  less  chemical  energy  than  do  the 
acids  themselves,  so  that  the  reactions  by  which  these  salts  are 
formed  are  exothermic.  One  consideration  forces  itself  upon  us 
when  we  study  the  structural  formula  of  fluosilicic  acid,  and  that 
is  the  impossibility  of  maintaining  the  theory  of  the  constant 
monovalence  of  fluorine  ;  for,  unless  we  wish  to  take  the  untenable 
position  that  fluosilicic  acid  is  a  compound  formed  of  finished 
molecules  of  silicon  fluoride  and  hydrofluoric  acid  (is  a  so-called 
"  molecular  compound  "),  we  must  look  upon  the  fluorine  atoms  as 
being  divalent  in  fluosilicic  acid  :  — 


Even  if  we  call  this  acid  a  molecular  compound  (SiF4,  2  HF),  such 
a  supposition  does  not  help  matters  in  the  least  ;  for  then  we  must 
regard  Si  F4  and  HF  as  still  having  chemical  affinity  at  their  dis- 
posal, which  supposition  is  contrary  to  the  theory  that  silicon  is 
only  tetravalent  and  fluorine  only  monovalent,  for  then  neither  of 
these  compounds  should  be  capable  of  further  union  after  Si  F4  and 
HF  have  been  formed.  So  long,  therefore,  as  our  present  theories 
of  valence  are  maintained,  we  must  regard  fluorine  as  being  both 
uni-  and  bivalent.  This  conclusion  is  strengthened  by  the  fact 
that  the  specific  gravity  of  hydrofluoric  acid  shows  that  substance 
to  have  a  molecule  corresponding  to  the  formula  H2F2. 

Fluosilicic  acid  is  known  only  in  solution  ;  when  evaporated  be- 
yond a  certain  concentration  it  breaks  down  into  silicon  tetrafluo- 
ride  and  hydrofluoric  acid,  much  as  ordinary  silicic  acid  does  into 
silicon  dioxide  and  water.*  A  similar  change  takes  place  with  the 
fluosilicates  ;  for  these  salts,  when  heated,  break  down  into  silicon 

*  A  hydrate  of  fluosilicic  acid,  H2  SiF6  ,  2  H2  O,  is  known;  it  is  a  crystal- 
line body  which  melts  at  19°  and  which  is  extremely  hygroscopic. 


304  SILICON   DIOXIDE. 

tetrafluoride  and  the  fluoride  of  the  metal  entering  into  the  salt 
formation  :  -  Si  F6  K2  =  Si  F4  +  2  KF. 

Almost  all  of  the  silicofluorides  are  soluble  in  water  ;  the  fluosilicate 
of  potassium  is,  however,  nearly  insoluble  ;  and,  as  almost  all  potas- 
sium salts  are  dissolved  by  water,  it  is  evident  that  fluosilicic  acid 
is  a  very  welcome  reagent  for  the  detection  of  potassium  compounds 
in  solution. 

Silicon  forms  but  one  well-known  oxide,  the  dioxide  Si  02  .  This 
substance  occurs  in  three  forms  :  crystalline,  cryptocrystalline, 
and  amorphous.  Crystallized  silicon  dioxide  is  dimorphous,  being 
found  as  quartz  *  and  as  tridymite.t  The  quartz  crystals  are  often 
colored  more  or  less  by  impurities  ;  when  the  color  so  produced  is 
purple  or  bluish  violet  the  crystal  is  called  amethyst.  The  crystal- 
lized variety  of  quartz  frequently  occurs  in  large  masses,  displaying 
no  crystalline  faces,  while  smaller  fragments  of  the  mineral  are 
found  as  a  constituent  of  the  granitic  rocks.  The  cryptocrystalline  $ 
varieties  of  quartz  show  the  greatest  diversity  of  color  and  appear- 
ance ;  they  generally  contain  more  or  less  water  and  are  more  read- 
ily acted  on  by  hydrofluoric  acid  than  the  crystallized  varieties. 
Examples  of  cryptocrystalline  quartz  are  chalcedony,  carnelian, 
agate,  onyx,  and  flint.  Sea  sand  consists,  for  the  most  part,  of 
quartz  finely  ground  by  the  action  of  the  water. 

Amorphous  silicon  dioxide  can  be  prepared  by  the  addition  of  an 
acid  to  a  soluble  silicate  :  — 


i  03  +  2  H  Cl  =  H2  Si  03  +  2  Na  Cl  ; 
Sodium  silicate.  Silicic  acid. 

and  by  then  heating  the  silicic  acid  until  all  water  is  expelled  ;  the 
silicon  dioxide  formed  in  this  way  is  a  white,  impalpable  powder 
which  can  be  readily  dissolved  in  alkalies  ;  when  heated  to  a  high 
white  heat  this  variety  of  the  dioxide  becomes  crystalline  and  can 
then  no  longer  be  dissolved  by  cold  alkalies. 

The  two  simplest  theoretical  hydrates  of  silicon  dioxide  are  ortho- 
silicic  acid,  Si  (0  H)4,  and  meta-silicic  acid,  Si  03  H2;  neither 

*  Hexagonal,  tetartohedral  ;  combinations  of  pyramid  and  prism. 

t  Asymmetric. 

\  Varieties  of  crystalline  minerals  in  which  the  crystals  are  so  small  as 
not  to  be  detected  by  the  eye  are  called  cryptocrystalline.  Such  rocks  are  fre- 
quently erroneously  termed  amorphous. 


DIALYSIS.  305 

of  these  acids  is  known  with  certainty,  but  a  solution  which  prob- 
ably contains  ortho-silicic  acid  can  be  obtained  by  the  following 
means:  by  adding  cold  dilute  hydrochloric  acid  to  a  very  dilute 
cold  solution  of  sodium  silicate  the  following  reaction  presumably 
takes  place  :  — 

i  03  +  2  H  Cl  =  H2  Si  03  +  2  Na  Cl  ; 


the  meta-silicic  acid  so  formed  then  unites  with  water  to  form  ortho- 
silicic  acid  :  — 

H2  Si  03  +  H2  0  =  H4  Si  04  . 

Under  the  conditions  of  the  reaction  there  is  no  separation  of  insol- 
uble silicic  acid,  as  there  is  when  the  solutions  are  more  concen- 
trated, or'when  they  are  heated.  If  the  clear  liquid  containing 
sodium  chloride  and  silicic  acid  is  put  in  a  vessel  the  bottom  of 
which  is  formed  of  a  membrane  such  as  parchment,  and  this  vessel 
is  then  placed  in  pure  water,  the  sodium  chloride  and  excess  of  hy- 
drochloric acid  will  pass  out  into  the  water,  while  the  silicic  acid 
will  remain  behind  in  the  solution.  The  process  by  which  this  sep- 
aration takes  place  is  osmosis,  and  the  silicic  acid  is  said  to  be  sepa- 
rated by  dialysis;  substances  which  are  able  to  pass  through  such  a 
membrane  are  called  crystalloids  ;  those  which  cannot  pass  through, 
colloids.  As  the  process  of  osmosis  is  one  of  extreme  importance 
in  animal  and  plant  life,  a  brief  discussion  of  some  of  the  principal 
facts  which  have  been  learned  regarding  it  may  not  be  out  of  place 
here.* 

When  a  layer  of  water  is  carefully  poured  over  any  aqueous  so- 
lution, the  two  liquids  will  not  remain  in  this  condition,  for  diffu- 
sion will  take  place  just  as  it  does  between  layers  of  different  gases 
(see  page  34),  so  that  the  solution  will  begin  to  rise  in  a  direction 
contrary  to  the  force  of  gravity  and  will  finally  completely  mix  with 
the  pure  water,  the  motion  only  ceasing  when  the  substance  in  solu- 
tion is  uniformly  distributed  throughout  the  mass  of  water.  This 
motion  can  be  arrested  by  placing  a  septum  between  the  water  and 
the  solution  ;  and  if  this  septum  is  of  such  a  material  as  to  allow 
water  to  pass  through,  but  not  the  dissolved  substance,  and  if,  fur- 
thermore, the  septum  is  in  the  shape  of  a  cell  which  can  be  covered 
by  an  air-tight  cap  which  is  so  constructed  that  it  can  be  connected 

*  See  Ostwald,  Outlines  of  General  Chemistry  (Walker),  for  a  more  com- 
plete description  of  this  topic. 


306  OSMOTIC    PRESSURE. 

with  a  manometer  by  means  of  a  glass  tube,  an  increase  of  pressure 
will  be  observed  in  the  interior  of  the  cell,  because  the  water  will 
force  its  way  in,  while  the  substance  in  solution  cannot  escape. 
Now,  a  remarkable  fact  is  observed  in  regard  to  this  pressure,  for  if 
the  temperature  is  kept  constant  the  pressure  will  be  proportional 
to  the  strength  of  the  solution;  thus,  with  a  solution  of  sugar,  the  fol- 
lowing pressures  were  observed,  — 

A  1  per  cent  solution  gave  a  pressure  of  535  m.  m. 
n  2    "      "  u  "    "         "        "  1016  "    " 

u  4    tt      u  a  a    a         a        a  2082  "    " 

These  observed  pressures  in  millimeters  are  nearly  in  the  pro- 
portion of  1:2:4.  This  law  of  osmotic  pressure,  which  is  true  of 
all  dissolved  substances,  is  exactly  like  that  regulating  the  pressure 
of  gases,  for  these  are  also  proportional  to  the  densities  (i.e.,  con- 
centrations). It  has  further  been  observed  that  temperature  has 
the  same  influence  on  osmotic  pressure  as  it  has  on  the  pressure  of 
gases ;  for  the  pressure  increases  proportionally  to  the  absolute  tem- 
perature, and  in  the  same  ratio  for  all  dissolved  substances.  The 
increase  in  pressure  for  each  degree  of  temperature  is  5|  5,*  a  frac- 
tion identical  with  that  obtained  as  an  increase  for  each  degree  in  the 
pressure  of  gases  which  are  kept  at  constant  volume  ;  this  relation  for 
osmotic  pressure  may  therefore  be  expressed  in  the  same  way  as  it  is 
for  gases.  If  we  know  the  osmotic  pressure  (  P0 )  at  0°,  then  at  t°  it 
will  be  P0  +  P0 .00367  t  =  P0  ( 1  -f  .00367  t)  (see  pages  172,  173). 
The  osmotic  pressure  of  a  substance  in  solution  has  the  same  value 
as  the  pressure  that  substance  would  exert  were  it  a  gas  occupying 
the  same  volume  as  the  solution.  It  seems  reasonable  to  suppose, 
therefore,  as  the  laws  which  govern  the  pressures  of  gases  also  hold 
good  for  osmotic  pressure,  that  the  substances  which  are  contained 
in  solution  are  present  in  such  solution  in  the  same  condition  as  that 
in  which  they  occur  in  gases,  i.e.,  as  the  individual  molecules. 

The  membranes  through  which  the  various  fluids  in  living 
organisms  must  find  their  way  by  osmosis  act  on  the  same  princi- 
ple as  the  septa  which  are  artificially  prepared ;  and,  as  the  same 
increase  of  temperature  causes  a  like  increase  of  osmotic  pressure 
in  all  fluids,  it  follows  that  solutions  which  are  in  osmotic  equilib- 
rium between  the  contents  of  a  living  cell  and  the  liquid  with- 

*  Exactly  .00367. 


ORTHO-   AND   META   SILICATES.  307 

out  at  any  given  temperature,  say  0°,  are  also  in  equilibrium  at 
38°. 

When  the  solution  containing  dialized  silicic  acid  is  evaporated, 
the  acid  congeals  to  a  gelatinous  mass,  which  is  then  no  longer 
soluble  in  water ;  and  when  this  is  separated  and  dried  the  remain- 
ing amorphous  powder  has  the  formula,  approximately,  of  H2  Si  03 ; 
this,  when  heated,  loses  water  and  forms  amorphous  silicon  dioxide. 

The  silicates  are  either  orthosilicates,  M4  Si  04 ,  or  metasilicates, 
M2  Si  03 ,  or  they  are  derived  from  more  complicated  silicic  acids 
which,  according  to  the  number  of  silicon  atoms  in  their  formula 
weights,  are  called  di-,  tri-,  or,  in  general,  polysilicates.  All  of  these 
salts  have  numerous  representatives  in  the  mineral  deposits  of  the 
earth.  The  formulae  of  a  few  of  these  are  given  in  the  following 
tabular  statement :  — 
Orthosilicates :  f  Olivin?  Mgg  g.  ^  ^  g.  ^  § 

I  Garnet,  Ca3Fe2  (Si  04)3.* 

Acid,    Si  J  |  Mica;  the  various  forms  of  this  mineral  are  complicated 

orthosilicates. 

f  Wollastonite,  Ca  Si  O3 . 
Leucite,  KA1  (SiO3)2. 
Beryll,  Be3Al2(Si03)6. 
Related  to  wollastonite,  but  of  more  complicated  struc- 


Metasilicates : 


J° 
i  OJ 


Acids,  Si  4  OH 
I  OH 


ture,  are  the  important  minerals,  hornblende  and 
augite. 

Disilicic  acid  is  formed  by  the  separation  of  one  molecule  of 
water  from  two  formula  weights  of  metasilicic  acid,  just  as  disul- 
phuric  acid  is  derived  from  two  of  sulphuric  (  page  150). 

(0  0)  (0  O) 

SU  OH-f  HO  V-Si  =  Si^  OH      HO  V- Si. 

(OH       HO)  (      -0-      ) 

Only  two  or  three  examples  of  disilicates  are  known.  Trisilicic 
acid  is  formed  by  separating  four  molecules  of  water  from  three 
formula  weights  of  Si  (OH  )4 ;  3  Si  (OH  )4  —  4  H2  O  =  Si3  O8  H4 . 

*  Of  the  twelve  hydrogen  atoms  in  three  formula  weights  of  orthosilic 
acid,  six  are  replaced  by  trivalent  iron  and  six  by  bivalent  calcium.  In  this 
silicate,  and  in  others,  two  or  more  formula  weights  of  the  acid  are  united  by 
an  atom  of  a  polyvalent  element,  which  replaces  hydrogen  atoms  belonging  in 
part  to  one,  and  in  part  to  another  formula  weight  of  the  acid  ;  a  simple 
example  of  such  a  case  we  have  encountered  in  the  formula  of  the  tertiary 
phosphate  of  calcium,  Ca3  (PO4)2,  (see  page  229). 


308 


TRISILICATES. 


Trisilicates: 

(O 

Si]  OH 

f  0 

(OH    „ 

Acid, 

Si  JOH 

'  O 

SijOH 

Orthoclase  (feldspar),  KAlSi3O8. 

Oligoclase  (soda,  lime  feldspar),  NaAl(Si3O8),  CaAl 
(AlSi2)08. 

(In  the  important  group  of  minerals  known  as  feldspars,  it 
not  infrequently  happens  that  a  portion  of  the  sili- 
con is  replaced  by  aluminium;  this  is  seen  in  the 
formula  of  oligoclase. ) 

The  quantitative  composition  of  the  silicates  shows  that  every 
one  of  them  can  be  considered  as  derived  from  one  of  the  above 
mentioned  acids.  A  number  of  basic  and  acid  silicates  also  exist. 
Among  acid  silicates,  kaolin  (clay)  H2  AL>  (Si04)2  +  H20,  may  be 
mentioned. 

The  silicates  are  such  extremely  important  minerals,  their  com- 
position is  so  varied  and  their  distribution  so  far  reaching,  that  the 
study  of  their  structure  forms  one  of  the  most  important  branches 
of  modern  mineralogy.  All  silicates,  excepting  those  of  the  alkali 
metals,  are  insoluble  in  water. 


GERMANIUM.  309 


CHAPTER   XLIL 

GERMANIUM  AND   ITS   COMPOUNDS. 
Germanium;  symbol,  Ge  ;  atomic  weight,  72.3. 

THIS  element  was  discovered  in  the  year  1886,  by  the  German 
chemist,  Clemens  Winckler,  and  is  especially  interesting  from  the 
fact  that  it  is  one  of  the  elements  the  existence  of  which  was  pre- 
dicted before  its  discovery.  This  prediction  was  based  upon  the 
fact  that,  when  the  elements  were  arranged  in  the  order  of  their 
atomic  weights  (page  17),  an  unfilled  gap  appeared  to  exist  be- 
tween gallium  (atomic  weight  69)  and  arsenic  (atomic  weight  75), 
(see  third  table,  chapter  xlix.),  which  gap,  as  the  nature  of  the  then 
known  elements  showed,  should  be  filled  by  a  representative  of  the 
carbon  family  (see  table  of  the  periodic  system).  Germanium  was 
subsequently  discovered  in  a  silver  ore  which  was  formerly  con- 
founded with  silver  sulphide  (argentite)  and  which  has  the  formula 
3  Ag2  S,  Ge  S2  and  is  termed  argyrodite  ;  the  properties  of  the  new 
element  were  found  to  be  in  accord  with  the  predictions  in  regard 
to  its  nature.  The  isolation  of  germanium  is  a  very  complicated 
process. 

Germanium,  owing  to  its  higher  atomic  weight,  must  be  much 
more  metallic  in  its  nature  than  silicon  ;  and,  indeed,  this  difference 
in  its  character  is  shown  by  the  non-existence  of  a  hydrogen  com- 
pound of  the  element.  The  metal  has  a  brilliant  metallic  lustre, 
and  is  readily  formed  from  its  oxide  by  reduction  with  charcoal  ;  and 
like  its  fellows,  carbon  and  silicon,  it  crystallizes  in  crystals  be- 
longing to  the  regular  system.  Its  specific  gravity  is  5.47,  and  its 
melting  point  less  than  that  of  silver  (954°).  The  metal  is  neither 
malleable  nor  ductile  ;  it  is  quite  brittle  and  can  be  readily  pounded 
to  a  powder;  in  this  respect  it  resembles  arsenic.  When  heated 
to  a  high  heat  in  the  air,  the  metal,  after  fusing,  oxidizes  to  form 
the  dioxide  Ge  02  .  Germanium  does  not  dissolve  in  hydrochloric 
acid  ;  it  is  oxidised  to  Ge  O2  by  nitric  acid  or  aqua  regia.  Hot  and 
concentrated  sulphuric  acid  dissolves  it  to  form  the  sulphate,  while 
the  acid  is  itself  reduced  to  sulphur  dioxide  (see  page  136). 


810  GERMANIUM;   COMPOUNDS   OF. 

Germanium  combines  with  all  of  the  halogens  to  form  com- 
pounds of  the  general  formula  GeK4,  where  E  represents  an  atom 
of  any  halogen.  Germanium  tetrachloride,  formed  by  passing 
chlorine  over  heated  germanium,  is  a  colorless  liquid  which  very 
much  resembles  the  corresponding  chloride  of  silicon,  Si  C14 ;  it  is 
decomposed  by  water,  and  boils  at  86° ;  the  specific  gravity  of  its 
vapor  (between  300°  and  740°)  is  7.43,  this,  H2  =  2,  is  213.9  ;  the  cal- 
culated molecular  weight  for  Ge  C14  is  214,  (Ge  =  72.3,  4  01  =  141.8) ; 
it  follows  from  this  that  the  formula  of  the  chloride  of  germanium 
corresponds  to  those  of  the  chlorides  of  silicon  or  carbon,  and  that  the 
maximum  atomic  weight  of  germanium  is  72.3  (see  page  72).  As, 
before  the  discovery  of  germanium,  no  element  was  known  to  exist 
having  an  atomic  weight  between  that  of  gallium  (69)  and  arsenic 
( 75  ),  and  as  an  element  belonging  to  the  carbon  family  and  having 
an  atomic  weight  of  approximately  72  would  evidently  find  a  fitting 
place  in  the  system  obtained  by  arranging  the  elements  in  the  order 
of  their  atomic  weights,  therefore,  the  gravimetric  quantity  of  ger- 
manium ( 72.3  ),  which  unites  with  141.8  (or  4  X  35.45 )  parts  by 
weight  of  chlorine,  is  probably  the  correct  atomic  weight  of  the 
element  in  question.  If,  at  any  time,  a  compound  of  germanium 
should  be  discovered  which,  with  a  known  molecular  weight,  should 
contain  proportionally  less  than  72.3  parts  by  weight  of  the  ele- 
ment, then  the  atomic  weight  which  is  at  present  accepted  will  have 
to  be  abandoned.  A  compound  of  germanium  and  chlorine,  GeCl2, 
corresponding  to  stannous  chloride,  Sn  C12 ,  has  also  been  mentioned, 
but  its  properties  have  not,  as  yet,  been  thoroughly  investigated. 
The  other  halogen  compounds  of  the  element  need  not  be  described, 
although  the  existence  of  a  germanium-chloroform,  Ge  H  C13  (a 
compound  corresponding  to  chloroform,  CHC13,  and  to  silico-chlo- 
roform,  Si  H  C13),  should  be  emphasized  as  showing  the  relationship 
between  germanium  and  the  preceding  elements  of  this  family. 
Germanium-chloroform  is  a  colorless  liquid  which  boils  at  72°  and 
which  is  readily  decomposed  by  water. 

Germanium  forms  two  oxides,  Ge  02  and  Ge  0.  The  former  is 
produced  either  by  burning  the  powdered  element  in  a  stream  of 
oxygen,  by  oxidizing  it  with  nitric  acid,  or  by  decomposing  the 
chloride  with  water.  It  is  a  white  powder,  somewhat  soluble  in 
water,  the  solution  probably  containing  the  hydroxide  H2  Ge  0  3, 
corresponding  to  metasilicic  acid,  H2  Si  03 .  Germanium  dioxide 


GERMANIUM;  COMPOUNDS  OF.  311 

acts  as  a  weak  acidic  anhydride,  dissolving  in  the  hydroxides  of  the 
alkali  metals  ;  it  has  no  basic  properties.  The  second  oxide,  Ge  0, 
is  an  unstable  substance  which  oxidizes  in  the  air,  has  a  hydroxide, 
Ge  (OH)2,  derived  from  it,  and  is  weakly  basic  in  its  character. 
It  is  a  powerful  reducing  agent. 

Two  sulphides  of  germanium,  Ge  S2  and  Ge  S,  are  known. 
These  correspond  to  the  sulphides  of  tin.  The  disulphide,  Ge  S2 , 
dissolves  in  the  sulphides  of  the  alkali  metals  to  form  sulpho-salts ; 
it  therefore  has  the  character  of  an  acidic  anhydride,  and  resembles 
the  sulphides  of  arsenic,  antimony,  tin,  and  carbon. 


312  TIN;  OCCURRENCE,  PREPARATION. 


CHAPTER   XLIII. 

TIN   AND  ITS   COMPOUNDS. 

Tin ;  symbol,   Sn ;  atomic  weight,  119. 

TIN,  the  third  element  of  the  carbon  family,  has  its  metallic 
properties  so  decidedly  pronounced  that,  disregarding  its  many 
resemblances  to  the  not-metals,  it  is  generally  classed  with  the 
metals.  In  reality  it  bears  about  the  same  relation  to  metals  and 
not-metals  as  antimony  does ;  however,  both  of  its  oxides  (stannous 
oxide,  Sn  0,  and  stannic  oxide,  Sn  02 )  have  basic  properties,  and, 
furthermore,  tin  cannot  form  a  hydrogen  compound ;  on  the  other 
hand,  both  oxides,  when  brought  in  contact  with  strong  bases,  can 
act  like  acidic  anhydrides. 

The  time  of  the  discovery  of  tin  is  not  known.  A  knowledge 
of  the  metal  has  been  attributed  to  the  Hebrews,  Greeks,  and  Phoe- 
nicians, but  no  certainty  exists  as  to  this.  Undoubtedly,  Pliny 
distinctly  mentions  tin  under  the  name  of  plumbum  candidum,  and 
moreover,  the  metal  was  used  by  the  Romans  for  covering  iron  in 
order  to  k^ep  that  metal  from  rusting.  The  term  stannum  dates 
from  the  fourth  century. 

Tin  is  one  of  the  comparatively  rare  elements,  and  its  occurrence 
in  the  free  state  as  a  mineral  is  somewhat  doubtful.  It  is  chiefly 
found  as  the  crystallized  dioxide,  Sn  02 ;  this  substance  (which  is 
termed  cassiterite  or  tinstone),  is  deposited  in  crystals  and  in  the 
massive  form  in  veins  traversing  granite,  gneiss,  and  mica  schist  in 
Cornwall  and  Devonshire,  the  Malay  peninsula,  New  South  Wales, 
and  Queensland.  Some  tin  has  also  been  discovered  in  the  United 
States.  An  impure  sulphide  of  tin  (stannite),  SnS2,  occasionally 
appears  in  mineral  deposits. 

The  tin  of  commerce  is  exclusively  prepared  from  tinstone.  The 
mineral  is  crushed  and  washed  and  then  heated  with  charcoal  ac- 
cording to  the  usual  metallurgical  process  (see  page  287).  The 
tin  which  melts  and  is  collected  at  the  bottom  of  the  furnace  is  gen- 
erally quite  impure,  for  it  contains  copper,  iron,  arsenic,  antimony, 
and  lead.  These  foreign  substances  are,  for  the  most  part,  removed 


TIN;  PROPERTIES.  313 

by  heating  the  crude  tin  to  a  temperature  jfist  above  its  melting 
point,  and  then  allowing  the  pure  metal  to  run  off  from  its  higher 
melting  alloys. 

Tin  is  nearly  silver-white,  with  a  metallic  lustre ;  it  is  scarcely 
corroded  when  exposed  to  the  air ;  it  is  soft  and  can  be  hammered 
and  rolled  into  thin  sheets  (tin  foil) ;  its  specific  gravity  is  7.3 ;  it 
melts  at  228°,  and  evaporates  at  a  temperature  between  1600°  and 
1700°.  The  metal  has  a  great  tendency  toward  crystallization, 
either  when  it  is  separated  from  its  compounds,  *  or  when  it  congeals 
after  fusion.  Tin,  like  carbon,  is  dimorphous,  occurring  both  in 
tetragonal  and  in  rhombic  crystals.  If  block  tin  is  cooled  to  a 
very  low  temperature,  or  even  if  it  is  allowed  to  stand  for  a  long 
time,  it  changes  into  a  grayish  powder ;  this  powder  will  reassume 
a  metallic  appearance  only  upon  being  fused.  This  amorphous 
form  of  tin  corresponds  to  amorphous  carbon  or  silicon.  Ordinary 
tin,  cast  into  forms,  assumes  a  crystalline  structure;  if  a  stick 
formed  of  the  metal  is  bent,  a  peculiar  crepitation  is  felt,  and  if  the 
operation  is  rapidly  repeated  several  times,  the  piece  of  tin  will 
become  quite  hot  at  the  place  of  bending ;  both  the  noise  and  the 
heat  are  caused  by  the  friction  of  the  minute  crystals,  one  upon  the 
other. 

Tin  is  attacked  by  acids  with  considerable  ease;  hydrochloric 
acid  dissolves  it  to  form  stannous  chloride :  — 

Sn-f  2HCl=SnCL,+2H, 

and  in  this  way  the  element  shows  its  metallic  nature ;  that  this  is 
not  very  pronounced,  however,  is  shown  by  the  fact  that  the  re- 
action between  tin  and  hydrochloric  acid  takes  place  much  more 
slowly  than  it  does  between  the  same  -acid  and  iron  or  zinc.  Hot 
and  concentrated  sulphuric  acid  dissolves  tin  to  form  stannous  sul- 
phate, while  the  acid  is  reduced  to  sulphur  dioxide  (see  page  137  ). 

*  By  electrolysis  of  stannous  chloride,  or  by  placing  a  piece  of  zinc  in  a 
solution  of  stannous  chloride,  the  zinc  then  takes  the  place  of  tin  in  the  salt:  — 

Zn  +  Sn  C12  =  Zn  C12  +  Sn. 

Such  substitutions  of  one  metal  for  another  in  salts  are  quite  frequently  met 
with,  but  are  not  surprising  if  we  compare  salts  with  acids,  for,  as  we  know, 
zinc  can  readily  replace  hydrogen  in  hydrogen  chloride,  and  why  not  tin 
in  stannous  chloride?  The  only  essential  is  that  stannous  chloride  should  have 
more  chemical  energy  than  zinc  chloride,  and  that  heat  should  be  given  off  in 
the  reaction. 


314  STANNOUS   CHLORIDE. 

Cold  and  dilute  nitric  acid  dissolves  tin  without  any  evolution  of 
gas  ;  the  tin  forms  stannous  nitrate,  while  the  nitric  acid  is  reduced 
to  ammonia  (page  206  [a]  )  ;  this  reaction  is  made  clear  by  the  fol- 
lowing equation  :  — 

4  Sn  +  10  HN03  =  4  Sn  (  N03  )2  +  NH4  N03  +  3  H2  0.* 

On  the  other  hand,  hot  and  concentrated  nitric  acid  oxidizes  tin  to 
insoluble  metastannic  acid,  H2  Sn  03  ,  and  is  itself  reduced  to  the 
lower  oxides  of  nitrogen  (page  207  [&]  ).f  Tin  shows  its  relation- 
ship to  the  not-metals  by  dissolving  in  alkaline  hydroxides  to  form 
salts  of  stannic  acid,  M2  Sn  03  . 

Tin  forms  two  series  of  compounds  ivith  the  halogens  ;  the  first 
of  these,  with  the  general  formula  SnX2  (where  X  represents  any 
halogen),  can  be  formed,  as  are  the  salts  of  other  metals,  by  dissolv- 
ing the  corresponding  oxide  in  halhydric  acids  :  — 

Sn  0  +  2  HX  =  SnX2  +  H20, 

while  the  compounds  SnX2  can  be  converted  into  those  of  the 
second  series,  Sn  X4  ,  by  the  addition  of  the  corresponding  halogen. 
The  two  chlorides  of  tin,  Sn  C12  ,  stannous  chloride,  and  Sn  C14  , 
stannic  chloride  (see  page  26),  are  the  most  important  of  these 
halogen  compounds. 

Stannous  chloride,  Sn  C12  ,  can  be  formed  by  dissolving  tin  or 
stannous  oxide,  Sn  0,  in  hydrochloric  acid  ;  when  anhydrous,  it  is  a 
crystalline  substance  which  melts  at  250°  and  boils  at  606°.  Its 
vapor  density  was  formerly  supposed  to  correspond  to  a  molecular 
weight  calculated  from  the  formula  Sn2  C14  ;  but  later  investigations 
have  shown  that  no  definite  specific  gravity  can  be  assigned  to  it. 
Stannous  chloride  dissolves  in  small  quantities  of  water  without 
change  ;  an  excess  of  the  solvent,  however,  partially  converts  it 
into  an  insoluble  basic  chloride  (see  page  253)  :  — 


Stannous  chloride  is  a  powerful  reducing  agent  ;  when  exposed  to 
the  air  it  absorbs  oxygen  and  changes  into  a  mixture  of  stannic 
chloride  and  the  basic  chloride  just  mentioned.  Stannous  chloride 
instantly  reduces  mercuric  chloride  to  mercurous  chloride  :  — 

*  Hydroxylanim  is  also  produced  by  this  reduction,  especially  if  hydro- 
chloric acid  is  present. 

t  Compare  with  antimony,  page  249. 


STANNIC   CHLORIDE.  315 

2  Hg  C12  +  Sn  C12  =  Sn  C14  +  2  Hg  Cl,* 

and  the  mercurous  chloride,  by  further  action,  is  even  finally 
changed  to  mercury.  Ferric  chloride  is  reduced  to  ferrous  chloride 
by  stannous  chloride  :  — 

2  Fe  C13  +  Sn  C12  =  Sn  C14  +  2  Fe  C12  , 

while  arsenic  trioxide  is  reduced  to  metallic  arsenic  by  the  same 
substance,  the  compound  of  tin  being,  in  this  case,  oxidized  to  stan- 
nic acid. 

When  alkaline  hydroxides  are  added  to  a  solution  of  stannous 
chloride,  insoluble  stannous  hydroxide  is  at  first  precipitated  :  — 

Sn  C12  +  2  KOH  =  Sn  (OH  )2  +  2  K  Cl, 

but  this  substance,  because  it  presents  in  a  slight  degree  the  charac- 
ter of  an  acid,  is  dissolved  by  an  excess  of  the  alkaline  solution  to 
form  stannites  which  are  salts  of  a  stannous  acid  having  the  formula 
H2  Sn2  03  ,  formed,  as  are  a  number  of  acids  which  we  have  already 
discussed,  by  the  separation  of  one  molecule  of  water  from  two 
formula  weights  of  the  hydroxide  :  — 

g     (OH       Sn-OH 


OH        Sn-OH 

The  potassium  compound,  formed  by  dissolving  stannous  hydrox- 
ide in  an  excess  of  potassium  hydroxide,  therefore,  has  the  formula 
K2  Sn2  03  ;  this  salt,  when  heated,  breaks  down  into  tin  and  potassium 
stannate  :  — 

Sn2  03  K2  =  Sn  03  K2  -f  Sn.t 

Stannic  chloride  can  be  formed  from  stannous  chloride  by  heating 
the  latter  substance  and  then  passing  dry  chlorine  over  it.  It  is  a 
colorless  liquid  which  fumes  in  the  air,  and  which  boils  at  114°;  it 
greedily  absorbs  moisture,  and  then  produces  crystals  having  the 
formula  SnCl4-f3H20.  Stannic  chloride  forms  a  series  of 

*  For  this  reason  stannous  chloride  is  used  as  a  reagent  for  soluble  salts 
of  mercury,  for,  as  mercurous  chloride  is  insoluble  in  water,  a  precipitate  of 
the  latter  is  formed  when  stannous  chloride  is  added  to  a  solution  containing 
mercury. 

t  See  pages  139  and  155.  This  reaction  is  similar  to  previous  ones  which 
we  have  studied. 


316  STANNOUS   OXIDE. 

double  salts  with  the  chlorides  of  the  alkali  metals,*  these  double 
salts  have  the  general  formula  SnCl6M2,  where  M  represents  an 
atom  of  an  alkali  metal ;  they  therefore  correspond  to  the  silico- 
fluorides,  Si  F6  M2  (see  page  302 ),  which  we  look  upon  as  being 
salts  of  fluosilicic  acid  (  H2  Si  F6 )  ;  there  is  consequently  no  reason 
why  the  conclusions  regarding  the  nature  of  fluorine  in  the  fluosili- 
cates  should  not  be  equally  applicable  to  chlorine  in  these  com- 
pounds of  tin;  indeed,  a  substance  H2  SnCl6  +  6  H2  0,  which  must 
be  considered  as  the  acid  from  which  these  double  salts  are  de- 
rived, has  in  all  probability  been  isolated. 

Tin  forms  two  oxides,  the  monoxide,  SnO,  and  the  dioxide, 
Sn  02 ;  these  correspond  to  the  oxides  of  carbon,  CO  and  C02 ;  the 
former  of  these  oxides  is  almost  altogether  basic  in  its  character, 
while  the  latter  most  frequently  acts  as  the  anhydride  of  an  acid. 

Stannous  oxide  can  best  be  prepared  by  heating  the  correspond- 
ing hydroxide  without  access  of  air.  It  is  a  dark  brown  substance 
which  dissolves  in  acids  to  form  stannous  salts,  or  in  alkalies  to 
form  stannites  ;  this  latter  reaction  has  already  been  fully  described 
under  stannous  chloride.  The  stannous  salts  are  colorless  when 
formed  from  a  colorless  acid,  and  are  readily  oxidized  when  in  con- 
tact with  the  air ;  those  insoluble  in  water  are  nearly  all  soluble  in 
dilute  hydrochloric  acid,  and  show  a  tendency  to  change  into  basic 
salts  on  the  addition  of  water. 

When  tin  is  heated  to  a  sufficiently  high  temperature  in  air  or 
in  oxygen,  it  burns  to  form  the  dioxide,  Sn  02 ;  this  substance  when 
•cold  is  a  white  powder,  but  when  hot  assumes  a  yellowish  color; 
after  being  exposed  to  a  high  temperature  for  some  time  it  becomes 
insoluble  both  in  acids  and  alkalies.  The  crystallized  variety  of 
the  oxide,  found  as  the  mineral  tinstone,  is  also  insoluble ;  the  only 
means  by  which  this  substance  can  be  brought  into  solution  is  by 
fusion  with  potassium  or  sodium  hydroxide,  when  the  respective 
stannates,  M2  Sn  03 ,  of  the  metals  are  formed. 

Two  stannic  acids,  identical  in  gravimetric  composition,  but  dif- 
fering in  physical  and  chemical  properties,  are  derived  from  the 
anhydride  Sn02;  both  have  the  formula  H2Sn03.t  Ordinary 

*  Stannous  chloride  also  forms  double  halides  with  the  chlorides  of  the 
alkali  metals ;  the  double  halides  of  potassium  are  K  Sn  C13  4-  H2  O  and  K2  Sn 
C14  +  2  H2  O . 

t  Two  orthostannic  acids  (Sn  (OH  )4  )  have  also  been  described.  See 
Neumann;  Monatshefte  fur  Chemie;  12,  515. 


STANNIC   ACIDS.  317 

stannic  acid  can  be  formed  by  adding  exactly  enough  potassium 
hydroxide  solution  to  a  solution  of  stannic  chloride  to  precipitate 
stannic  hydroxide  :  — 

Sn  C14  +  4  KOH  =  Sn  (OH)4  +  4  K  Cl ; 

this  substance  separates  as  a  jelly  which  resembles  silicic  acid. 
When  this  is  dried  it  loses  water  and  changes  to  a  gum-arabic  like 
mass  which  approximately  has  the  formula  H2  Sn  03 .  This  variety 
of  stannic  acid  is  readily  soluble  both  in  acids  and  alkalies.  The 
other  form  of  stannic  acid,  generally  called  metastannic  acid,  is 
presumably  a  polymeric  form  of  ordinary  stannic  acid,  so  that  if 
the  molecule  of  the  latter  had  the  formula  H2  Sn  03 ,  that  of  the 
former  would  be  expressed  by  n  (H2  Sn  03 ).  A  more  correct  system 
of  nomenclature  is  one  by  which  ordinary  stannic  acid  is  termed 
a  stannic  acid,  while  metastannic  acid  is  called  ft  stannic  acid. 
ft  stannic  acid  is  produced  in  the  form  of  an  insoluble  white  powder 
when  tin  is  oxidized  by  means  of  strong  and  hot  nitric  acid ;  when 
carefully  dried  in  a  vacuum  it  has  the  formula  H2  Sn  03 ;  it  is 
insoluble  in  acids,  and  when  heated  to  redness,  loses  water  and 
changes  into  the  dioxide  of  tin.  If  ft  stannic  acid  is  digested  with 
hydrochloric  acid  for  some  time,  the  hydrochloric  acid  then  poured 
off  and  pure  water  added,  the  stannic  chloride  so  formed  will  dis- 
solve ;  on  addition  of  alkalies  to  this  solution,  however,  ft  stannic 
acid  is  once  more  precipitated.  When  ft  stannic  acid  is  boiled  with 
sodium  hydroxide  it  is  converted  into  the  ft  stannate  of  sodium, 
which  can  be  dissolved  by  pouring  off  the  excess  of  caustic  soda 
solution  and  then  adding  pure  water.  Fusion  with  solid  caustic 
alkalies  converts  ft  stannic  acid  into  salts  of  ordinary  stannic  acid. 
Stannic  chloride,  or  what  amounts  to  the  same  thing,  the  solu- 
tion of  either  stannic  acid  or  ft  stannic  acid  in  hydrochloric  acid, 
resembles  the  chlorides  of  the  not-metals  in.  so  far  as  it  is  converted 
into  the  corresponding  acid  by  boiling  with  water ;  this  decompo- 
sition is,  however,  only  partial,  for  if  a  solution  of  ordinary  stan- 
nic acid  in  hydrochloric  acid  is  boiled  in  a  retort,  the  volatile 
stannic  chloride  passes  over  unchanged  in  company  with  the  water 
and  hydrochloric  acid,  while  very  little  stannic  acid  will  remain 
behind ;  on  the  other  hand,  a  solution  of  stannic  chloride  derived 
from  ft  stannic  acid  is  completely  decomposed  into  that  acid  by 
boiling,  while  no  stannic  chloride  whatever  will,  pass  over. 


318  TIN;  SULPHIDES. 

The  stannates  of  the  alkali  metals  are  the  only  ones  which  are 
soluble  in  water ;  in  that  way  these  salts  resemble  those  of  silicic 
and  carbonic  acids. 

Tin  forms  two  sulphides,  with  the  formulae  Sn  S  and  Sn  S2 ;  in 
structure  these  compounds  correspond  to  the  oxides.  The  monosul- 
phide,  Sn  S,  can  be  produced  by  adding  hydrogen  sulphide  to  an 
acidulated  solution  of  a  stannous  salt.  The  sulphide  is  a  brownish 
black  powder,  which  is  insoluble  in  dilute  acids,  but  is  dissolved  by 
concentrated  hydrochloric  acid  or  by  aqua  regia ;  in  the  latter  case 
stannic  chloride  is  formed.  Simple  sulphides  of  the  alkalies  scarcely 
attack  it ;  it  is  dissolved  by  the  polysulphides  (see  page  155,  foot- 
note), because  the  latter  sulphurize  it  to  form  salts  of  sulphostannic 
acid,  in  a  manner  exactly  similar  to  their  action  on  the  trisulphide 
of  antimony  (see  page  256).  The  disulphide  of  tin,  Sn  S2 ,  can  be 
precipitated  from  a  weakly  acid  solution  of  stannnic  chloride  by 
means  of  hydrogen  sulphide ;  it  forms  a  yellow  precipitate  which  is 
not  dissolved  by  dilute  acids,  but  which  is  soluble  in  strong  hydro- 
chloric acid  or  aqua  regia ;  with  the  latter  reagent  it  forms  stannic 
chloride.  As  would  be  expected,  owing  to  the  weakly  metallic  na- 
ture of  tin,  the  disulphide  acts  as  if  it  were  an  acidic  anhydride ; 
it  is,  therefore,  readily  attacked  by  either  the  hydroxides  or  sul- 
phides of  the  alkalies  (see  page  256  and  foot-note),  for  in  the  former 
case  a  mixture  of  stannate  and  sulphostannate  is  formed,*  while  in  the 
latter  the  sulphostannate  alone  is  produced,  f  This  behavior  of  tin 
is  very  much  like  that  of  antimony  or  arsenic  under  similar  circum- 
stances (see  pages  245  and  255).  On  addition  of  acids  to  the  solu- 
tion of  sulphostannates,  stannic  sulphide  is  precipitated,  Sn  S3  K2  -f 
2  H.  Cl  =  2  K  Cl  +  H2  S  +  Sn  S2 ,  for  the  sulphostannic  acid  which 
would  be  formed  at  once  breaks  down  into  hydrogen  sulphide  and 
the  disulphide  of  tin.  (It  will  be  noticed  that  the  formulae  of  the 
sulphostannates  correspond  to  those  of  the  salts  of  dithio-sulpho- 
carbonic  acid  [see  page  293].) 

The  compounds  of  tin  show  an  almost  perfect  concordance  with 
those  of  carbon  and  silicon,  when  the  formulae  alone  are  considered  ; 
chemically,  however,  the  acid  nature  of  the  substances  in  question 
is  materially  reduced  because  of  the  metallic  character  belonging 
to  tin.  This  metallic  character  becomes  much  more  pronounced  in 

*  3  Sn  S2  +  6  KOH  =  Sn  O3  K^  +  2  Sn  S3  K2  +  3  H.2  O. 
t  SnS2 


TIN  AND   CARBON  ;    COMPARISON   OF. 


319 


the  next  element  of  the  family  (lead),  so  that  the  oxides  of  that 
element  are  for  the  most  part  basic. 

The  relationship  between  the  compounds  discussed  in  this  chapter 
will  be  seen  from  the  following  table  :  — 


TIN    AND    CARBON. 


OXIDES. 

CHLORIDES. 

ACIDS. 

SULPHIDES. 

CO 
CO,' 

SnO 
SnO2 

CC14 

SnCl2 
SnCl4 

H2Sn2O3 

H2SnO3 

SnS 

Sn  S2 

Forms  no  sulpho  salts. 
Forms  sulpho  salts,  M2  X  S3. 

H2CO3 

CS2 

The  oxide  Sn  O  is  both  basic  and  acidic. 

It  dissolves  in  acids  as  follows  :   Sn  O  +  2  HX  =  Sn  X2  +  H2  O. 

It  dissolves  in  bases  as  follows  :   2  Sn  O  +  2  MOH  =  Sn2  O3  M2  +  H2  O. 

The  oxide  Sn  O2  is  both  basic  and  acidic. 

It  dissolves  in  halhydric  acids  as  follows  :  — 

SnO.,  +  4HX  =  SnX4  +  2H20. 
It  dissolves  in  bases  as  follows  :  — 

Sn  02  +  2  MOH  =  Sn  O3M2  +  H2  O. 


a  STANNIC  ACID,  H2  Sn  O3. 


STANNIC  ACID,  H2  Sn  O3. 


Derived  from  an  oxide  soluble  in  acids 
and  alkalies. 

The  chloride,  SnCl4,  derived  from  it  is 
volatile  in  the  vapors  of  dilute  hydrochloric 
acid. 


Sodium  hydroxide  readily  dissolves  it  in 
the  cold,  forming  ordinary  stannate  of 
sodium. 

Ordinary  stannic  chloride,  on  addition  of 
alkalies,  precipitates  ordinary  stannic  acid. 


Derived  from  an  oxide  insoluble  in  acids 
and  alkalies. 

The  chloride,  SnCl4,  derived  from  it  is 
not  volatile  with  water  vapors  (metastan- 
nic  chloride);  it  decomposes  into  hydrochlo- 
ric acid  and  ft  stannic  acid  when  the  solution 
is  boiled. 

Sodium  hydroxide,  when  boiling,  forms 
ft  stannate  of  sodium,  which  is  soluble  in 
water. 

ft  stannic  chloride,  on  addition  of  alkalies, 
precipitates  ft  stannic  acid. 


320  LEAD;    OCCURRENCE,   PREPARATION. 


CHAPTER   XLIY. 

vj 

LEAD    AND    ITS    COMPOUNDS. 

Lead  ;  symbol,  Pb;  atomic  weight,  206.95. 

THE  element  is  seldom  found  as  the  native  metal.  Its  most 
important  natural  compound  is  the  sulphide,  Pb  S,  which  occurs 
widely  distributed  as  galena  (also  termed  galenite).  The  carbon- 
ate (cerassite,  PbC03),  the  sulphate  (anglesite,  PbS04),  the  chro- 
mate,  phosphate,  and  molybdate  are  also  not  infrequently  met 
with. 

Lead  is  one  of  the  metals  which  has  been  known  since  the  old- 
est times,  having  been  familiar  to  the  Israelites.  The  Eomans 
made  much  the  same  use  of  the  metal  as  we  do  at  the  present  time, 
for  they  constructed  water-pipes  of  it,  and  prepared  a  solder  com- 
posed of  two  parts  of  lead  and  one  of  tin. 

The  lead  which  is  met  with  as  a  commercial  product  usually 
contains  copper,  iron,  and  traces  of  silver.     It  is  prepared  by  heat- 
ing the  sulphide  with  finely  divided  iron :  — 
Pb  S  +  Fe  =  Fe  S  +  Pb. 

This  operation  is  conducted  in  tall  furnaces,  which,  in  shape,  re- 
semble the  blast  furnaces  for  the  manufacture  of  pig  iron.  Another 
method  for  the  production  of  the  metal  consists  in  roasting  the 
sulphide  in  a  current  of  air,  by  which  means  it  is  in  part  oxidized, 
so  that  a  mixture  of  the  sulphate,  oxide,  and  sulphide  are  formed, 
and  when  this  mixture  is  heated  to  a  higher  temperature  the  sul- 
phate and  oxide  are  finally  reduced  by  the  sulphide,  which  is  still 
present;  the  sulphur  passes  off  in  the  form  of  sulphur  dioxide, 
while  the  lead  remains  behind.  The  crude  lead  contains  silver, 
antimony,  arsenic,  copper,  iron,  and  zinc;  the  oxidizable  impurities 
are  removed  by  melting  the  metal  in  the  air.  When  silver  is  pres- 
ent in  sufficient  quantity  to  pay  for  its  isolation,  the  entire  mass  is 
melted  and  subjected  to  a  blast  of  air,  by  which  means  lead  oxide 
is  produced ;  the  latter  melts  and  is  run  off  from  the  surface,  while 
the  silver  remains  behind  unchanged.  The  lead  oxide  so  formed 


LEAD  ;    PROPERTIES.  321 

can  then  be  once  more  reduced  to  lead  by  means  of  charcoal. 
Another  method  for  removing  the  silver  consists  in  melting  the 
silver-bearing  lead  and  then  allowing  the  mass  to  cool  slowly;  pure 
lead  crystallizes  at  first ;  this  can  be  removed  by  means  of  a  ladle, 
while  the  molten  mass  remaining,  which  is  very  rich  in  silver,  can 
be  treated  according  to  the  method  mentioned  above ;  or  the  silver 
can  be  removed  by  melting  the  lead  with  zinc,  for  zinc  mixes  with 
lead  in  a  small  proportion  only,  while  it  is  able  to  dissolve  all  of  the 
silver. 

Lead  is  a  metal  with  a  bluish-gray  color  and  metallic  lustre; 
it  is  malleable  and  easily  fused,  its  melting  point  is  330°,  and  it 
can  be  boiled  at  a  high  white  heat.*  The  metal  crystallizes  in 
octahedra ;  its  specific  gravity  is  11.4 ;  when  freshly  cut  it  has  a 
bright,  metallic  surface,  which,  however,  soon  becomes  covered  with 
a  layer  of  the  oxide,  and  this  protects  the  remainder  from  further 
corrosion.  If  a  piece  of  zinc  is  placed  in  a  solution  of  a  lead  salt, 
the  lead  will  separate  in  a  crystalline  form,  while  the  zinc  takes  its 
place  :  — 

.       Zn  +  Pb  (N08)2  =  Zn  (X03)2  +  Pb; 

similar  substitutions  are  not  infrequently  met  with  in  the  chemis- 
try of  other  metals  (see  page  313  and  foot-note). 

When  lead  is  covered  with  water  which  is  in  contact  with  the 
air,  it  becomes  covered  with  a  layer  of  lead  hydroxide ;  the  latter 
substance  is  to  a  certain  extent  soluble  in  water ;  as  a  consequence 
water  which  has  passed  through  new  lead  pipes  contains  more  or 
less  of  the  hydroxide  in  solution,  and  may,  for  this  reason,  prove 
to  be  highly  poisonous.  Hard  water  gradually  changes  the  hy- 
droxide into  the  entirely  insoluble  carbonate,  so  that,  in  time,  the 
pipes  become  covered  with  a  protective  coating.  In  dealing  with 
lead  pipes,  however,  care  must  be  taken  to  have  no  decaying  organic 
substances  present,  for  such  impurities  may  remove  the  carbonate 
and  greatly  increase  the  solubility  of  the  lead. 

Lead  is  not  readily  attacked  by  hydrochloric  f  or  cold  sulphuric 
acid ;  hot  and  concentrated  sulphuric  acid  has  some  effect  on  it,  as 

*  The  boiling  point  lies  between  1450°  and  1600°  (Carnelley  and  Williams; 
Journ.  Chem.  Soc. ;  35,  563). 

t  Hot  hydrochloric  acid  attacks  lead  to  a  considerable  extent  (Sharpies; 
€hem.  News;  50,  126). 


. 

322  LEAD   MONOXIDE. 

is  evinced  by  the  fact  that  commercial  sulphuric  acid  always  con- 
tains lead ;  diluted  nitric  acid  readily  dissolves  the  metal  to  form 
lead  nitrate.  Concentrated  nitric  acid  has  but  little  effect. 

Lead  enters  quite  readily  into  a  number  of  alloys,  some  of 
which  have  already  been  mentioned  (see  page  251)  ;  the  metal  is 
easily  amalgamated  by  mercury. 

Lead  forms  four  oxides,  Pb2  0,  suboxide  of  lead ;  Pb  0,  lead 
monoxide ;  Pb2  O3 ,  lead  trioxide  ;  and  Pb  02 ,  lead  dioxide,  or  lead 
hyperoxide.  Of  these,  the  oxides  PbO  and  Pb  02  are  the  most 
important;  the  monoxide,  PbO,  corresponds  to  carbon  monoxide, 
and  the  dioxide,  Pb  02 ,  to  carbon  dioxide. 

The  oxide  of  lead,  Pb  0,  can  easily  be  produced  by  heating  the 
nitrate  (see  page  201),  or,  like  other  oxides  of  weakly  pronounced 
metals,  it  can  be  formed  by  heating  the  hydroxide :  — 

Pb(OH)2  =  PbO  +  H20. 

This  oxide  of  lead  is  ordinarily  a  yellow  powder,  which  is  easily 
melted  to  an  orange-colored  mass  (litharge) ;  but  another,  red, 
modification  of  the  oxide  can  be  obtained  by  heating  lead  hydrox- 
ide to  150°.  The  hydroxide  of  lead  separates  as  a  white  precipi- 
tate when  a  base  is  added  to  a  solution  of  a  lead  salt :  — 

Pb  ( N03 )2  +  2  Na  OH  =  Pb  (OH )2  +  2  Xa  NOS, 

and,  like  stannous  hydroxide,  it  is  both  a  base  and  an  acid.  As  a 
result  of  its  basic  properties,  it  readily  dissolves  in  acids  to  form 
salts :  — 

Pb  (OH)2  +  2  HN03  =  Pb  (N03)2  +  2  H2  0  ; 

and  because  of  its  acid  properties  it  dissolves  in  pronounced 
alkalies :  — 

Pb  (OH)2  +  2  KOH  =  Pb  (OK)2  +  2  H2  0; 

but  the  hot  saturated  solution  deposits  lead  oxide  on  cooling  (com- 
pare with  stannous  hydroxide,  page  315).  The  salts  of  lead  can 
be  formed  by  dissolving  either  the  oxide  or  hydroxide  in  acids,  and, 
being  salts  of  a  pronounced  metal,  they  are  more  or  less  stable; 
they  are  poisonous  when  in  a  soluble  form.  Among  the  most  im- 
portant salts  of  lead  are  the  chloride,  PbCl2,  sulphateJPb  S04 ,  and 
chromate,  Pb  Cr  04 ,  which  are  insoluble  or  nearly  insoluble  in 
water;  they  can  therefore  be  produced  from  the  soluble  lead  salts 
by  the  addition  of  a  soluble  chloride,  sulphate,  or  chromate.  The 


LEAD   DIOXIDE.  323 

soluble  carbonates  of  the  alkalies  when  added  to  the  solution  of  a 
lead  salt  cause  a  precipitate  of  basic  carbonate  of  lead,  Pb(OH)2 
Pb  C03 ;  this  substance  is  white  lead.  Among  the  important  soluble 
lead  salts,  the  acetate  (Pb(C,  H80,)2,  sugar  of  lead)  may  be  men- 
tioned. One  interesting  fact  as  regards  lead  salts  is  the  isomor- 
phism which  the  sulphate  displays  with  the  sulphates  of  the  very 
pronouncedly  metallic  alkaline  earths  (calcium,  barium,  and  stron- 
tium), and  the  isomorphism  of  the  carbonate  with  the  carbonates 
forming  the  arragonite  group.*  This  relationship  shows  us  that 
the  pronounced  metallic  properties  of  lead  have  caused  it  to  depart 
so  far  from  the  family  type  that  its  salts  resemble  those  of  the  most 
characteristic  divalent  metals ;  this  isomorphism  is  also  displayed  in 
the  case  of  some  other  lead  compounds. 

The  oxide  of  lead  next  in  importance  to  the  monoxide,  Pb  0,  is 
the  dioxide  or  hyperoxide,  Pb  02 ;  this  substance  has  only  very  weakly 
basic  or  acid  properties.  It  belongs  to  the  class  of  hyperoxides  of 
which  manganese  dioxide  is,  perhaps,  the  best  known  representa- 
tive. The  hyperoxides  are  all  neutral,  or  nearly  neutral  bodies, 
which,  when  heated  with  sulphuric  acid,  give  off  oxygen  and 
change  to  the  sulphate  derived  from  the.  oxide  MO  as  a  base,  and 
which,  when  treated  with  hydrochloric  acid,  liberate  chlorine.  The 
dioxide  of  lead  is  occasionally  found  in  nature ;  it  can  be  formed 
in  the  laboratory  by  treating  the  oxide  Pb304f  with  nitric  acid, 
or  by  oxidizing  the  acetate  of  lead  with  a  solution  of  chloride  of 
lime. 

Lead  dioxide  is  a  dark  brown  powder,  which  is  a  powerful 
oxidizing  agent ;  indeed,  it  can  oxidize  sulphur  dioxide  so  readily 
that  the  heat  of  the  reaction  may  even  cause  it  to  glow,  provided  it 
is  finely  divided  and  placed  in  an  atmosphere  of  the  gas.  When 
in  contact  with  strong  bases  it  dissolves  to  form  salts  of  an  acid, 
plumbic  acid,  H2  Pb  03 ,  which,  in  formula,  .corresponds  to  carbonic 
acid,  but  which,  like  the  latter,  does  not  exist  in  the  free  state ; 
only  a  very  few  salts  of  this  acid  are  known.  On  the  other  hand, 

*  The  carbonates  of  calcium,  barium,  manganese,  and  iron. 
t  The  red  oxide  of  lead,  Pb3  O4,  may  be  looked  upon  as  a  mixture  of  the 
oxides  Pb  O  and  Pb  O2:  — 

2PbO  +  PbO2  =  Pb3O4; 

the  nitric  acid  dissolves  out  the  monoxide  Pb  O  and  leaves  the  dioxide  Pb  O2; 
compare  this  formula  with  Mn3  O4  and  Fe3  O4 . 


324  MINIUM.      LEAD   SULPHIDE. 

the  dioxide  is  soluble  in  some  acids,  although  the  salts  which  are 
presumably  formed  by  this  action  have  not  been  isolated. 

The  red  oxide  of  lead  (  Pb3_04 ,  so-called  minium)  is  produced  by 
carefully  oxidizing  finely  powdered  litharge  at  300°  to  400°  —  at 
a  higher  temperature  it  breaks  down  with  the  liberation  of  oxygen 
and  the  regeneration  of  litharge.  This  oxide  is  extensively  used  as 
a  pigment  under  the  name  of  red  lead.  The  suboxide  of  lead,  Pb2  0, 
and  the  sesquioxide,  Pb2  O3 ,  are  of  little  importance. 

The  metallic  character  of  lead  is  so  predominant  that  the  nature 
of  its  oxides  is  very  much  at  variance  with  that  of  the  oxides  of 
the  elements  at  the  beginning  of  this  family  ;  indeed,  in  the  forma- 
tion and  character  of  the  compounds  Pb3  04  and  Pb  02 ,  lead  very 
much  resembles  manganese,  while  the  isomorphism  of  its  salts  with 
those  of  calcium,  barium,  and  strontium,  brings  it  in  close  connection 
with  the  alkaline  earths ;  on  the  other  hand,  it  is  like  carbon,  sili- 
con, and  tin,  for  its  oxide  Pb  02  shows  weakly  acid  properties.  In 
fact,  the  chemical  characteristics  of  lead  are  not  very  marked  in  any 
direction,  nor,  indeed,  is  this  neutral  behavior  unexpected,  for  we 
find  it  to  be  quite  a  general  fact  that  the  elements  with  high  atomic 
weights  and  specific  gravities  *  display  no  very  pronounced  chemical 
properties ;  the  crowding  of  a  large  mass  into  a  small  space,  as  is 
the  case  with  these  elements,  is  therefore  unfavorable  for  the 
manifestation  of  striking  chemical  phenomena. 

Lead  forms  only  one  sulphide  which  has  been  accurately  stud- 
ied, the  monosulphide  Pb  S.  This  substance  is  found  as  the  min- 
eral galena,  crystallized  in  cubes  of  metallic  appearance.  In  the 
laboratory  it  can  be  produced  either  by  direct  combination  of  the 
elements,  or  by  precipitation  from  acid  solutions  of  lead  salts  by 
means  of  hydrogen  sulphide ;  when  so  prepared  it  is  a  black,  amor- 
phous powder  which  is  not  attacked  by  cold  hydrochloric  acid,  but 
which  is  attacked  by  that  substance  when  it  is  hot  and  concen- 
trated. Oxidizing  agents,  such  as  nitric  acid,  change  it  into  the 
insoluble  sulphate  of  lead,  and  a  similar  transformation  is  brought 
about  by  roasting  in  the  air. 

In  the  following  table  the  most  important  compounds  which 
have  representatives  in  the  chemistry  of  a  number  of  elements  of 
this  family  are  placed  side  by  side;  those  acids  which  are  not 

*  I.e.,  with  small  atomic  volumes  (see  chapter  xlix.). 


COMPOUNDS  OF  CARBON  FAMILY. 


325 


known   in  a  free  state,   but  salts  of  which  exist,  are  placed  in 
parentheses  :  — 


OXIDES.                                                                                                ACIDS. 

Monoxides. 
Dioxides. 

CO,  ,  Ge  O,  Sn  O,  Pb  O.              ,  ,  ,  H2  Sn2  O3)  Pb  (OH)2. 
C02,  Si  O2,  Ge  O2,  Sn  O2)  Pb  O2.    (H2  CO3),  H2  SiO3,—  ,H2  Sn  O3)(H2PbO3) 

SALTS. 

M2  CO3,  M2  Si  O3,  ,  M2  Su  O3,  M2  Pb  O3. 

CHLOBIDES. 

SULPHIDES. 

Dichlorides. 
Tetrachlorides. 

,  ,GeCl2,SnCl2, 
Pb  C12. 
C  C14,     Si  C14,     Ge  C14, 
SnCl4,  PbCl4.* 

Monosulphides. 
Bisulphides. 

CS,      ,    ,    SnS, 
PbS. 
CS2,  SiS2,  GeS2,  SnS2, 

Presumably  formed  because  the  dioxide  dissolves  in  acids. 


326  ELEMENTS   OF   BORON   FAMILY. 


CHAPTER   XLV. 

THE   ELEMENTS   OF   THE    BORON   FAMILY    (THE   EARTHS). 

THE  elements  of  the  family  of  which  boron  is  the  member  with 
the  smallest  atomic  weight,  are  the  last  which  will  be  considered 
where  any  purely  not-metallic  element  occurs.  As  has  been  re- 
peatedly mentioned,  the  groups  of  elements  become,  as  a  whole, 
more  metallic  in  their  nature  as  the  atomic  weights  diminish ;  this 
fact  is  readily  recognized  by  comparing  the  various  families  in  the 
order  in  which  they  have  been  studied,  as  is  shown  in  the  table  on 
page  266.  In  the  nitrogen  group  there  are  four  elements,  nitrogen, 
phosphorus,  arsenic,  and  antimony,  which  could  be  classed  with  the 
not-metals ;  in  that  of  carbon  there  are  but  two,  carbon  and  silicon ; 
while  in  the  one  under  consideration,  boron  alone  appears  to  us  with 
pronouncedly  not-metallic  characteristics,  while  even  this  element 
is  unable  to  form  a  gaseous  hydrogen  compound  of  sufficient  sta- 
bility to  render  an  accurate  study  of  its  properties  practicable. 

The  elements  comprising  the  boron  family  are  given  in  column 
1,  those  of  the  carbon  family  are,  for  purposes  of  comparison,  placed 
in  column  2. 

1.     Boron,  atomic  weight,  11  2.     Carbon,  atomic  weight,  12 
Aluminium,  "        "        27  Silicon,  "         "        28.4 

Gallium,        "        "       69  Germanium, "         "        72.3 

Indium,         "       "      113.7  Tin,  "         "       119. 

Thallium,      "        "     204.18  Lead,  "         "      206.95 

The  highest  valence  toward  oyxgen  displayed  by  the  elements 
of  the  nitrogen  family  is  five,  as  is  shown  by  the  existence  of  the 
pentoxides,  X2  05 ;  in  the  carbon  family  the  power  of  uniting  with 
oxygen  is  exhausted  when  the  dioxide,  X0.2,  in  which  the  element 
is  quadrivalent,  is  reached  ;  while  in  the  boron  group,  with  a  return 
to  the  type  of  oxide  shown  by  nitrogen  and  its  fellows,  the  highest 
valence  displayed  toward  oxygen  is  only  three,  so  that  the  character- 
istic oxides  of  this  group  have  the  formula  X2  03 .  These  oxides, 
of  course,  suffer  a  diminution  in  their  acidic  character  the  greater 
the  atomic  weight  of  the  element  forming  them  is  ;  so  that  boron 
trioxide,  B2  03 ,  acts  as  an  acidic  anhydride  under  all  circumstances ; 


ELEMENTS    OF   BOKON   FAMILY;    COMPARISON.  327 

aluminium  trioxide,  A12  03 ,  is  both  basic  and  acidic ;  the  oxide  of 
gallium,  Ga203,  displays  the  same  character;  while  the  oxide  of 
indium,  In2  03 ,  is  almost  exclusively  basic,  for  it  with  difficulty 
dissolves  in  caustic  alkalies,  and  the  unstable  compound  so  formed 
is  broken  down  by  warming  the  solution ;  finally,  the  trioxide  of  thal- 
lium, T12  03 ,  is  not  affected  by  the  reagents  in  question  ;  from  the 
above  comparative  statement  it  follows  that  the  oxides  X2  O3  are 
more  basic  the  greater  the  atomic  weight  of  X.  As  the  elements  in 
the  family  increase  in  atomic  weight,  they  display  the  same  ten- 
dency to  form  a  number  of  oxides  which  is  observed  in  the  case  of 
lead  in  the  carbon  family  ;  this  fact  will  become  apparent  by  a 
study  of  the  following  table :  — 

Boron  forms  one  oxide,  B.2  O3 ; 

Aluminium  forms  one  oxide,  A10  O3 ; 

Gallium  forms  two  oxides,  Ga  O  and  Ga2  O3 ; 

Indium  forms  two  oxides,  InO  and  In2  Og; 

Thallium  forms  three  oxides,  T12  O,  T12  O3 ,  and  Tl  O2 . 

The  oxides  with  least  amount  of  oxygen,  derived  from  any  given 
element  in  this  family,  capable  of  forming  more  than  one  oxide  are, 
without  exception,  basic  in  their  character,  while  the  trioxides,  with 
the  exception  of  that  of  thallium,  are  both  basic  and  acidic.  The 
majority  of  the  salts  which  contain  an  element  which  is  a  member 
of  this  family  are  derived  from  the  trioxide  X2  03 . 

What  is  true  of  the  oxides  is  also  true  of  the  halogen  com- 
pounds ;  the  elements  with  high  atomic  weights  each  are  capable  of 
forming  more  than  one  chloride,  bromide,  or  iodide,  while  alumin- 
ium and  boron  are  confined  to  one  apiece.  The  trihalide,  like  the 
trioxide,  is  the  compound  common  to  all  of  the  members  of  the 
family.  The  relationship  between  these  compounds  can  be  seen  by 
examining  the  following  table ;  as  will  be  noticed,  the  rule  which 
holds  good  in  all  the  preceding  families,  namely,  that  the  boiling 
points  of  the  trichlorides  are  higher  the  greater  the  molecular 
weight  of  the  compound,  is  also  without  exception  in  this  group :  — 

BC13,  liquid,  boils  at  17°; 

Al  C13,  solid,  melts  at  190°,  boils  at  183°; 

Ga  C13 ,  solid,  melts  at  75°.5,  boils  at  220°; 

In  C13 ,  solid,  volatilizes  at  red  heat  without  melting; 

Tl  C13 ,  gives  off  chlorine  when  heated,  and  changes  to  the  chloride  Tl  Cl. 

In  addition  to  the  chlorides  given  above,  gallium  forms  a  compound  with 
the  formula  Ga  C12 ,  indium  the  chlorides,  In  Cl  and  In  C12  ,  while  thallium  also 
has  a  monochloride,  Tl  Cl. 


328  BORON;    OCCURRENCE,   PREPARATION. 


CHAPTER  XLVI. 

BORON   AND    ITS    COMPOUNDS. 

Boron  y  symbol,  B  ;  atomic  weight,  11. 

IN  its  physical  characteristics  boron  bears  a  marked  resemblance 
to  carbon  (the  element  having  the  next  highest  atomic  weight  to  its 
own) ;  but  in  the  chemistry  of  its  oxides  and  chlorides,  boron  is 
very  much  like  the  members  of  the  nitrogen  family. 

Boron  is  never  found  uncombined  in  mineral  deposits.  Its  com- 
pounds, which  not  infrequently  occur  in  nature,  are  either  salts  of 
boric  acid  or  the  acid  itself.  The  most  important  of  these  minerals 

are:  — 

Borax  (tinkal)  NaH  (BO2)2  ,  4  H2O; 

Borocalcite,  CaH2  (BO2)4  ,  5  H2  O. 

The  borates  of  other  metals,  for  instance,  of  iron  and  magnesium,  are 
also  found,  while  solutions  of  boric  acid  sometimes  occur  in  lagoons  of  volcanic 
regions. 

Although  the  element  was  not  isolated  until  1807,*  and  was  not 
accurately  described  until  1824,t  its  compounds,  especially  tinkal, 
occurring  as  they  do  in  mineral  deposits,  were  known  in  very  early 
times,  the  natural  borax  having  become  familiar  to  Europeans  by 
importations  from  India. 

Boron  can  be  isolated  by.  heating  the  oxide  (  B2  03 )  with  sodium, 
or  potassium  fluoborate  (KBF4  )  with  potassium.  When  so  obtained 
it  is  an  amorphous,  brownish-black  powder  which  greatly  resembles 
silicon  in  appearance ;  it  is  quite  readily  dissolved  by  melted  alu- 
minium, and,  when  the  metal  containing  the  boron  is  cooled,  the  ele- 
ment separates  in  the  form  of  reddish-yellow,  diamond-like  crystals, 
which  are  very  hard  and  lustrous.  When  the  aluminium  has  been 
heated  to  a  temperature  only  just  above  its  melting  point,  then  the 
dissolved  boron  appears  in  a  graphitoidal  form.  Elementary  boron, 
therefore,  displays  modifications  similar  to  those  belonging  to  carbon. 
The  specific  gravity  of  boron  is  2.68 ;  it  is  infusible. 

*  By  Guy  Lussac  and  Thenard, 
t  By  Berzelius. 


BORON;    HALIDES.  329 

When  heated  in  the  air,  amorphous  boron  burns  to  form  the 
trioxide  B203;  the  same  modification  of  the  element  is  readily 
oxidized  by  nitric  acid,  or  even  by  concentrated  sulphuric  acid,* 
boric  acid,  B  03  H3  ,  being  produced.  Boron  can  unite  directly  with 
chlorine,  bromine,  with  some  metals,  and  with  nitrogen. 

A  gaseous  hydrogen  compound  of  boron  was  not  known  until 
quite  recently,  f  It  was  then  prepared  in  an  impure  state  by  treat- 
ing the  boride  of  magnesium  with  hydrochloric  acid  ;  this  method 
corresponds  to  the  one  by  which  silicon  hydride,  Si  H4  ,  and  pure 
arsine  and  stibine  are  produced.  Hydrogen  boride  is  very  unstable  ; 
it  burns  in  air  or  in  oxygen  with  a  bright  green  flame,  and  it  is 
slightly  soluble  in  water.  The  formula  assigned  to  the  gas  is  B  H3  , 
but  a  more  extended  investigation  of  its  composition  is  necessary. 

The  halogen  compounds  of  boron  correspond  to  the  general 
formula  of  B  X3  ,  so  that,  in  structure,  they  are  identical  with  the 
trihalides  of  the  nitrogen  family.  The  trichloride  and  trifluoride  are 
the  only  representatives  of  these  compounds  which  we  need  consider. 
The  trichloride  is  formed  by  the  direct  union  of  chlorine  and  boron  ;  $ 
when  first  discovered  it  was  supposed  to  be  a  gas  at  ordinary  tem- 
peratures, but  subsequent  investigations  proved  it  to  be  a  liquid 
with  a  boiling  point  at  17°.  Being  the  chloride  of  a  not-metal, 
boron  trichloride  is,  of  course,  readily  decomposed  by  water,  and 
normal  or  ortho  boric  acid  results  from  this  decomposition  :  — 

Cl       HOH  OH 


CI  +  HOH         OH 

It  will  be  noticed  that  this  change  is  parallel  to  the  one  undergone 
by  phosphorus  trichloride  when  in  the  presence  of  water  (see  pages 
81  and  220). 

Boron  trifluoride  is  a  colorless  gas,  and  is  interesting  because  its 
chemical  properties  are  much  like  those  of  the  tetrafluoride  of  sil- 
icon (see  page  302).  It  can  be  prepared  by  treating  the  dry 

*  Compare  the  action  of  amorphous  carbon  (charcoal)  on  nitric  acid 
(note  59  of  appendix)  and  on  sulphuric  acid  (page  136). 

t  In  1881,  by  Jones  and  Taylor.  See  also  Sabatier;  Comptes  Rendus;  112,  t 
865. 

t  By  passing  chlorine  over  an  intimate  mixture  of  boron  and  carbon  heated 
to  redness.  This  method  of  preparation  is  exactly  like  the  one  employed  in 
the  formation  of  the  chloride  of  silicon,  Si  C14  . 


330  BORIC   ACID. 

trioxide  of  boron  with  hydrofluoric  acid,*  so  that,  in  this  reaction, 
boron  trioxide  is  a  base  :  — 

B203  + 6HF=2BF3  +  3H20. 

Boron  trifluoride,  when  passed  into  water,  undergoes  a  decomposition 
similar  to  that  experienced  by  silicon  tetrafluoride,  for  it  breaks 
down  into  boric  and  fluoboric  acids  (see  page  303)  :  — 

4  BF3  +  3  H2  0  =  3  HBF4  +  B  (OH  )3 . 

Fluoboric  acid  bears  a  very  close  resemblance  to  fluosilicic  acid. 
It  forms  fluoborates  with  the  general  formula  of  MBF4,  when  it  is 
brought  in  contact  with  the  hydroxides  of  the  alkali  metals  ;  and  in 
these  compounds  fluorine  must  necessarily  be  considered  as  a  biva- 
lent element  for  reasons  identical  with  those  brought  forward  in  the 
discussion  of  silicofluorides  on  page  303. 

The  only  oxide  of  boron  is  the  trioxide,  B2  03 .  The  latter  can 
be  formed  either  by  burning  amorphous  boron,  or,  as  is  more  expe- 
dient, by  heating  the  hydroxide  B  (OH  )3  ( boric  acid)  to  redness ; 
the  trioxide  is  a  glass-like  mass  which  is  soluble  in  water  and  which, 
as  it  volatilizes  only  at  a  very  high  temperature,  will,  when  heated 
with  the  salts  of  other  volatile  acids,  finally  decompose  those  salts 
and  form  borates.  For  the  same  reason,  fused  boric  acid  is  able  to 
dissolve  the  great  majority  of  metallic  oxides.75 

Boric  acid  (  B  [OH  ]3 )  not  infrequently  occurs  in  natural  deposits, 
it  being  found  in  a  crystalline  state  in  the  neighborhood  of  the 
iumaroles  t  of  Tuscany ;  the  acid  so  found  is  called  "  sassolin." 
'The  water  of  the  lagoons  in  this  region  contains  about  one-tenth  per 
•cent  of  boric  acid ;  this  amount  is  increased  to  as  much  as  one  per 
'Cent  by  collecting  the  water  in  cisterns  and  then  allowing  the  vapors, 
•charged  with  boric  acid,  to  condense  in  these  receptacles  ;  the  cistern 
water  is  finally  evaporated  in  flat  leaden  pans,  which  are  warmed  by 
the  steam  which  is  escaping  from  the  earth ;  the  solid  residue  which 
remains,  containing  about  75  per  cent  of  boric  acid,  is  purified  by 

*  By  mixing  the  trioxide  Bg  Og  with  fluorspar,  fusing,  and  then  adding 
sulphuric  acid  to  the  mass  when  cool.  Compare  this  formation  with  that  of 
silicon  tetrafluoride. 

t  Fumaroles  are  jets  of  water  vapor  which  escape  from  fissures  in  the 
earth  in  volcanic  regions ;  these  vapors  condense  on  the  surface  and  form  small 
lagoons,  which  are  kept  boiling  by  the  continued  injection  of  hot  vapors. 


BOKATES.  331 

recrystallization.     Boric  acid  is  also  prepared  for  commercial  use  by 
decomposing  natural  borax  *  and  borocalcite  by  means  of  acids. 

Boric  acid  is  a  white,  crystalline,  flaky  solid  which,  to  the  touch, 
has  a  peculiar  fatty  feeling  ;  it  is  tolerably  soluble  in  water,  one 
part  of  the  acid  being  taken  up  by  twenty-six  parts  of  the  solvent 
at  ordinary  temperatures.  Boron  trioxide  is  also  quite  soluble  in 
alcohol;  when  this  solution  is  lighted,  the  solvent  burns  with  a 
characteristic  green  flame.76  f 

The  usual  form  of  boric  acid  is  orthoboric  acid  ;  when  this  sub- 
stance is  heated  to  100°,  water  is  given  off  and  metaboric  acid  is 
produced  :  — 

OH 


(OH 
B(OH)3=B02H  +  H20.* 

When  metaboric  acid  is  heated  to  160°  it  is  changed  to  tetraboric 
acid,  which  has  the  formula  H2  B4  07  ,  its  structure  being  on  a  plan 
similar  to  those  of  di-  and  trisilicic  acids  (see  pages  307  and  308). 
The  most  common  salt  of  boric  acid,  commercial  borax,  is  derived 
from  tetraboric  acid.  Finally,  when  tetraboric  acid  is  fused,  all  of 
the  hydroxyl  groups  are  separated  in  the  form  of  water,  and  the 
anhydride,  B203,  is  formed. 

The  metaborates,  MB02  ,  and  the  tetraborates,  M2  B4  07  ,  are 
the  most  stable  salts  of  boric  acid.  The  orthoborates,  M3  B  03  ,  are 
easily  decomposed  ;  the  organic  derivatives  of  orthoboric  acid  are, 
however,  tolerably  stable  substances. 

The  not-metallic  properties  of  the  elements  of  the  boron  family 
are  not  very  p'ronounced;  indeed,  the  element  with  next  higher 
atomic  weight,  namely,  aluminium,  is  almost  invariably  metallic  in 
its  nature  ;  so  that,  as  a  consequence,  we  would  expect  boron  triox- 
ide to  act  as  a  base  under  some  circumstances  ;  and  this  is  found  to 
be  the  case  when  we  consider  that  the  oxide  dissolves  in  hydrofluo- 
ric acid  to  form  boron  trifluoride,  BF3  ,  and  is  algo  made  apparent 
by  the  existence  of  a  phosphate  of  boron,  BPO4  . 

*  Found  in  large  quantities  in  some  alkaline  lakes,  for  instance  in  Borax 
Lake,  California. 

t  This  green  flame  is  due  to  the  formation  of  the  triethyl  ester  of  acid,  i.e.  ; 
boric  acid  in  which  the  three  hydrogen  atoms  have  been  replaced  by  ethyl. 

|  Compare  with  the  table  on  page  227. 


332  BORON   NITRIDE. 

Only  one  oxy-,  or  acid-chloride  of  boron,  having  a  character 
similar  to  the  chlorides  of  sulphur  and  of  phosphorus  (see  pages 
156  and  221)  exists.  This  substance  has  the  formula  B  0  Cl ;  it  is 
derived  from  metaboric  acid  by  replacing  the  hydroxyl  group  with 
chlorine  ;  it  is  decomposed  by  water. 

The  sulphide  of  boron  in  formula  corresponds  to  the  oxide ;  it 
can  be  formed  by  direct  union  of  the  elements.  No  sulpho-salts 
derived  from  this  compound  are  known;  indeed,  it  is  decomposed 
with  the  greatest  violence  when  brought  in  contact  with  water,  so 
that  in  the  latter  respect  it  resembles  the  sulphide  of  silicon. 

Boron  is  one  of  the  few  elements  that  is  capable  of  direct  union 
with  nitrogen.  The  nitride  of  boron,  BN,  *  is  a  white  solid,  formed 
by  heating  an  intimate  mixture  of  borax  and  ammonium  chloride. 
Like  cyanogen,  it  is  a  stable  substance ;  it  is  not  attacked  either  by 
acids  or  alkalies,  but  when  heated  in  a  current  of  steam  it  breaks 
down  into  boric  acid  and  ammonia :  -^ 

BN  +  3  H2  0  =  B  (OH  )3  +  NH3 . 

The  occurrence  of  boric  acid  in  the  fumaroles  of  Tuscany  is  attrib- 
uted to  the  subterranean  decomposition  of  the  nitride  of  boron  by 
means  of  water  vapor. 

*  This  nitride  is  therefore  analogous  to  cyanogen,  (CN)2  . 


ALUMINIUM  ;   OCCURRENCE.  333 


CHAPTER   XLVII. 

ALUMINIUM   AND   ITS    COMPOUNDS. 

Aluminium;  symbol,  Al;  atomic  weight,  27  ;  specific  gravity,  2.56. 

ALUMINIUM  never  occurs  as  the  uncombined  metal.  Its  oxide, 
hydroxides,  fluoride,  and  silicates  are,  however,  very  widely  dis- 
tributed in  mineral  deposits.  The  chief  aluminium  compounds 
which  are  found  in  the  form  of  mineral  individuals  are  :  — 

Coftindum;  A12  O3 ,  named  ruby  when  found  in  red,  transparent  crystals. 

Diaspor;  A1O2H,  this  hydroxide,  in  formula,  corresponds  to  metaboric 
acid. 

Beauxite;  A1.2  O  (OH)4 . 

Hydrargyllite ;  A1(OH)3,  this  substance  is  the  normal  aluminium  hy- 
droxide. 

Cryolite;  3NaF,AlF3. 

Spinell ;  Mg  A12  O4,  a  magnesium  salt  of  the  hydroxide  Al  O2  H . 

In  addition  to  these  oxides  and  hydroxides  and  the  salts  derived  from 
them,  aluminium  forms  a  large  number  of  silicates.  Among  the  orthosilicates 
which  contain  aluminium  are :  — 

Garnet;  Ca3  A12  (Si  O4)3 . 

Muscovite;  KH2  A13  (SiO4)3  ;  and  allied  to  muscovite  are  the  various 
micas. 

Kaolin  (clay);  H4Al2Si2O9. 

Some  of  the  most  important  meta-silicates,  for  example,  hornblende  and 
augite,  have  already  been  mentioned,  as  have  also  the  very  important  polysili- 
cates  which  belong  to  the  feldspar  group  (see  page  308). 

So  extended  is  the  distribution  of  aluminium  in  the  mineral 
kingdom  that  it  can  safely  be  asserted  -that  the  element  is  a 
constituent  of  the  greater  number  of  natural  silicates.  Basic 
sulphates  and  phosphates  of  aluminium  are  also,  not  infrequently, 
met  with. 

Although  its  compounds  are  so  widely  distributed,  aluminium 
itself  was  not  discovered  until  1827,  when  Wohler  prepared  the 
metal  by  heating  powdered  aluminium  chloride  with  potassium. 
At  a  later  date,  St.  Claire  Deville  introduced .  the  use  of  sodium, 
which  is  comparatively  cheap,  in  place  of  the  very  dear  metal 


334  ALUMINIUM;  PREPARATION. 

potassium  and,  at  a  still  later  date,  Kose  improved  the  process  by 
using  cryolite,  Al  F3 ,  3  KF,  instead  of  aluminium  chloride.  Until 
very  recently,  however,  all  of  the  aluminium  of  commerce  was 
prepared  by  heating  the  chloride  of  that  metal  with  sodium,  by 
which  reaction  sodium  chloride  and  aluminium  are  formed  :  — 

Al  C13  +  3  Na  =  Al  +  3  Na  Cl. 

In  the  last  four  years  this  expensive  process  has  given  way  to 
the  electrolytic  production  of  aluminium.*  The  material  operated 
on  is  commercially  pure  aluminium  oxide  (A12  03) ,  which  is  ob- 
tained by  chemical  methods  from  beauxite  or  other  aluminous 
minerals.  A  bath  of  melted  cryolite  is  maintained  in  a  fused 
condition  by  passing  an  electric  current  of  several  thousand 
amperes  through  carbon  cylinders  hung  in  the  melted  cryolite, 
which  is  itself  contained  in  an  iron  tank  lined  with  carbon.  The 
suspended  rods  constitute  the  anode  and  the  tank  the  cathode. 
When  pulverized  aluminium  oxide  is  stirred  into  this  bath  it  is 
promptly  dissolved,  and  then  becomes  the  electrolyte  (see  page  13), 
which  is  acted  on  by  the  current.  The  metallic  aluminium,  in  a 
melted  condition,  goes  to  the  cathode,  while  the  oxygen  which  is 
at  the  same  time  separated  from  the  decomposing  oxide,  burns  up 
the  carbon  anodes,  which,  in  consequence,  require  frequent  renewal. 
By  reason  of  this  improvement  in  the  manufacture  of  the  metal, 
the  price  of  aluminium  has  fallen  very  greatly  of  late  years. 

The  specific  gravity  of  aluminium  is  2.56,  t  its  melting  point  is 
700° ;  the  metal  cannot  be  volatilized.  Aluminium  is  a  good  con- 
ductor of  heat ;  it  conducts  electricity  almost  eight  times  as  well  as 
iron  does.  When  covered  with  concentrated  nitric  acid,  the  metal  is 
transferred  into  a  condition  in  which  it  is  not  further  attacked  by 
the  acid ;  when  the  metal  is  in  this  state  it  will  generate  an  electric 
current  when  placed  in  contact  with  ordinary  aluminium.  A  metal 
acting  in  this  way  is  said  to  be  in  the  "  passive  state."  No  adequate 
explanation  of  this  phenomenon  has  as  yet  been  given.  $  Alu- 
minium is  easily  attacked  by  hydrochloric  acid. 

*  Personal  communication  from  Prof.  J.  W.  Langley,  whom  I  wish  to 
take  this  opportunity  of  thanking  most  heartily.- 

t  This  specific  gravity  is  for  cast  aluminium ;  hammered  aluminium  has  a 
specific  gravity  of  2.67. 

J  It  is  supposed  that,  in  the  case  of  aluminium,  the  metal  becomes  covered 
with  a  layer  of  hydrogen  which  protects  it  from  further  attack.  Diluted  sul- 


ALUMINIUM,   PROPERTIES.  335 

Aluminium,  because  of  its  small  specific  gravity,  its  toughness, 
and  the  difficulty  with  which  it  is  attacked  by  the  corroding  agents 
which  ordinarily  come  in  contact  with  a  metal  in  general  use,  will, 
in  the  future,  have  its  commercial  usefulness  limited  only  by  the 
cost  of  its  production,  and,  as  we  have  seen,  the  latter  is  constantly 
diminishing.  A  number  of  aluminium  alloys  are  finding  extended 
application;  perhaps  the  most  important  of  these  is  aluminium 
bronze,  composed  of  about  ten  parts  of  aluminium  to  ninety  parts  of 
copper ;  this  composition  is  more  easily  worked  than  ordinary  bronze, 
is  tougher,  is  tarnished  with  difficulty,  and  has  the  color  of  gold. 
Perfectly  pure  aluminium  is  not  tarnished,  either  in  dry  or  moist 
air. 

Aluminium  forms  but  one  series  of  compounds  with  the  halo- 
gens;  these  compounds  have  the  general  formula  A1X3,  and  of 
these  the  chloride  (Al  C13)  is  the  most  important.  Aluminium  chlo- 
ride is  produced  by  heating  powdered  aluminium  in  a  current  of 
dry  chlorine.  Although  a  solution  of  aluminium  oxide  or  hy- 
droxide in  hydrochloric  acid  undoubtedly  contains  aluminium  chlo- 
ride, just  as  is  the  case  in  a  similar  solution  of  arsenic  trioxide 
(see  page  181),  nevertheless,  the  salt  cannot  be  isolated  by  evapo- 
rating the  liquid  because,  as  the  trichloride  is  the  halogen  com- 
pound of  a  metal  with  very  weakly  pronounced  metallic  properties, 
it  at  once  breaks  down  into  the  hydroxide  of  aluminium  and  hydro- 
chloric acid :  — 

(  Cl  +  HOH  f  OH 

Al  •]  Cl  +  HOH  =  Al  •]  OH  +  3  HC1. 

(  Cl  +  HOH  (  OH 

Aluminium  chloride  is  a  white,  crystalline  solid  which  fumes  when 
in  contact  with  the  air ;  it  greedily  absorbs  moisture,  and,  while 
giving  off  hydrochloric  acid,  changes  into  aluminium  hydroxide ;  it 
boils  at  180°.  The  vapor  density  of  aluminium  chloride  was 
formerly  supposed  to  be  9.34,  if  air  is  taken  as  unity ;  this  specific 
gravity  corresponds  to  a  molecular  weight  of  267,  and  a  formula 
A12  C16 .  This  experimental  evidence  inaugurated  a  theory  of  the 
quadrivalence  of  aluminium  in  its  trichloride,  for  the  structure  of 

phuric  acid  scarcely  attacks  aluminium  under  ordinary  atmospheric  pressure, 
but  if  the  metal  is  placed  in  a  vacuum,  then  bubbles  of  gas  are  given  off  and 
the  action  goes  on.  Similar  phenomena  are  observed  with  nitric  acid.  (See 
Ditte;  Comptes  Rendus;  110,  573.) 


336  ALUMINIUM;    HALIDES. 

the  latter  compound,  were  the  molecule  to  have  the  formula  A12  C16 
would  be  as  follows  :  — 

Cl  )  (  Cl 

Cl  V-  Al  —  Al  J  Cl 

Cl )  (  Cl. 

A  similar  constitution  was  assigned  to  the  trichloride  of  iron, 
which  substance  likewise  was  supposed  to  have  a  molecule  corre- 
sponding to  the  formula  Fe2Cl6.  The  latest  investigations  of 
Nilsson  and  Pettersson*  have  shown,  however,  that  the  vapor 
density  of  9.34  found  for  aluminium  chloride  is  only  incidental 
to  a  certain  temperature,!  and  this  specific  gravity  steadily  dimin- 
ishes as  the  heat  is  increased,  until  it  reaches  4.6  at  800° ;  this 
latter  number  remains  constant  up  to  1500°  ;  from  this  it  is  evident 
that  the  molecular  weight  of  aluminium  chloride  is  133.5;  the 
formula  of  this  compound  is  consequently  A1C13,$  so  that  alumin- 
ium must  be  regarded  as  a  trivalent  element.  Even  if  no  vapor 
density  determinations  of  the  chloride  of  iron  had  been  made,  it 
would  seem  probable  that  the  trichloride  has  the  formula  of  FeCl3, 
for  iron  can  replace  aluminium  in  isomorphous  mixtures,  so  that  a 
difference  in  the  valence  of  the  two  elements  in  compounds  derived 
from  the  trioxide  M2  03  seems  scarcely  probable.  It  is  unnecessary, 
however,  to  call  to  our  aid  any  such  reasoning,  because  the  final 
investigations  regarding  the  vapor  density  of  ferric  chloride  have 
definitely  decided  that  that  substance,  like  aluminium  chloride,  has 
a  molecule  represented  by  the  formula  Fe  C13 . 

The  halogen  compounds  of  aluminium  possess  the  power  in  a 
marked  degree  of  forming  double  salts  with  the  halides  of  other 
metals.  Formerly  these  double  salts  were  looked  upon  as  molec- 
ular additions,  formed  of  a  finished  molecule  of  some  halogen 
salt  of  aluminium  united  with  the  halide  of  an  alkali  metal. 
Such  a  theory  really  means  that  we  have  no  knowledge  of  the 
structure  of  such  compounds,  although,  in  maintaining  it  we  must 
believe  that  the  molecules  have  a  residuum  of  chemism  at  their  dis- 
posal. The  theory  of  simple  molecular  addition  would  no  longer 

*  Zeitschrift  fur  Physikalische  Chemie,  iv.,  206.  t  357°. 

I  This  discovery  is  further  borne  out  by  the  fact  that  certain  organic  deriv- 
atives of  aluminium,  which  have  been  obtained  as  gases,  undoubtedly  contain 
that  element  in  a  trivalent  form  (aluminium  triethyl,  aluminium  trimethyl, 
aluminium  acetylaceton). 


ALUMINIUM;    DOUBLE    SALTS.  337 

be  tenable  if  any  of  the  double  salts  could  be  obtained  as  gases 
with  unchanged  composition,  and  in  the  case  of  the  double  chloride 
of  sodium  and  aluminium,  Deville  states  that  the  compound  can  be 
vaporized  without  separating  it  into  molecules  of  Al  C13  and  Ka  Cl. 
A  theory  which  is  of  late  being  regarded  with  considerable  favor  is 
the  one  which  supposes  that,  in  the  double  halides,  the  halogen 
compounds  of  such  weakly  metallic  elements  as  arsenic,  antimony, 
bismuth,  or  aluminium  assume  the  role  of  acidic  anhydrides,  while 
the  halides  of  the  alkali  metals  are  the  bases  ;  and,  in  order  to 
maintain  such  a  theory,  the  assumption  is  inevitable  that  the  atoms 
of  the  halogens  can,  under  certain  circumstances,  become  divalent  ; 
such  a  belief  is  strengthened  by  the  existence  of  acids  like  fluosi- 
licic  acid,  H2  Si  F6  ,  and  fluoboric  acid,  HBF4  (see  pages  303  and  330). 
That  structures  similar  to  those  of  the  salts  of  oxy-acids  are  pos- 
sessed by  the  double  halide  salts  seems  more  than  probable  if  we 
consider  the  following:  "AYhen  a  halide  of  any  element  combines 
with  the  halide  of  an  alkali  metal  to  form  a  double  salt,  the  num- 
ber of  molecules  of  the  alkali  salt  which  are  added  to  one  molecule 
of  the  other  halide  is  never  greater,  and  is  generally  less,  than  the 
number  of  halogen  atoms  contained  in  the  latter."  *  It  must  be 
confessed,  however,  that  the  rule  has  a  number  of  exceptions.  The 
chlorides  of  aluminium  and  potassium,  and  of  aluminium  and 
sodium  have  the  formulae  A1C13,  KC1,  and  A1C13,  XaCl,  and  if  we 
regard  A1C13  as  analogous  to  an  acidic  anhydride,  and  KC1  and 
Nad  as  analogous  to  bases,  the  structure  of  these  compounds 
would  be  as  follows  :  — 

^Jv  and 
C12K 

the  parallelism  between  these  formulae  and  those  of  the  correspond- 
ing oxy-compounds  becomes  apparent  when  we  consider  the  structure 
of  the  latter  ;  namely  :  — 


The  double  fluorides  of  aluminium  are  of  two  kinds,  the  first  of 

(•  -p 
which,  with  the  general  formula  Al  -j  -p2iyp  are  constructed  similarly 

to  the  chlorine  compounds  ;  while,  second,  with  the  general  formula 

*  Remsen;  American  Chemical  Journal;  14,  85.    See  also  Remsen;  Chem- 
istry, p.  461. 


338  ALUMINIUM   OXIDE. 

(F2M  (OM* 

Al  J  F2  M,  correspond  to  the  theoretical  ortho-aluminates,  Al  J  OM 

(F2M  (OM. 

In  addition  to  the  substances  which  have  just  been  discovered,  a 
number  of  compounds,  in  jwhich  aluminium  chloride  is  a  base,  are 
known;  an  example  would  be  the  compound  A1C13,  PCl5.f  The 
existence  of  these  double  chlorides  once  more  reminds  us  of  the 
great  resemblance  between  the  chemistry  of  the  halogens  and  that 
of  oxygen  (see  pages  62,  63,  and  157  ). 

Aluminium  forms  but  one  oxide,  the  trioxide  A12  03 .  This  sub- 
stance occurs  as  the  mineral  corundum,  which,  when  finely  divided 
and  mixed  with  oxide  of  iron,  is  called  emery.  The  transparent, 
red  crystals  of  the  oxide  are  called  ruby.  Aluminium  trioxide  is 
produced  when  aluminium  is  heated  to  a  high  temperature  in  air  or 
in  oxygen,  or  when  the  hydroxide  is  heated ;  the  latter  substance, 
because  it  is  insoluble  in  water,  is  precipitated  from  solutions  of 
aluminium  salts  by  adding  ammonia  water  :  — 

AlCl3-t-3NH3  +  3H20  =  Al  (OH  )3  +  3  NH4  Cl ; 
and  when  heated  :  - 

2  Al  (OH  )3  =  A12  03  +  3  H2  0. 

The  oxide  which  has  been  heated  to  redness,  or  the  naturally  occur- 
ring crystalline  varieties,  are  insoluble  both  in  acids,  water,  or  solu- 
tions of  the  alkalies ;  they  can  be  brought  into  solution  by  fusing 
with  caustic  alkalies.  The  oxide  which  has  not  been  heated  is  both 
basic  and  acidic,  for  it  dissolves  in  acids  to  form  the  salts  of  alu- 
minium, and  in  bases  to  form  aluminates. 

Several  hydroxides  are  derived  from  aluminium  trioxide.  The 
first  of  these,  the  normal  hydroxide  Al  (OH  )3 ,  is  precipitated  from 
solutions  of  aluminium  salts  by  ammonia  water  or  caustic  alkalies ; 
in  using  the  latter  reagents,  however,  care  must  be  taken  not  to  add 
an  excess,  otherwise  solution  of  the  hydroxide  with  the  formation 
of  an  aluminate  takes  place. $  The  remaining  hydroxides  are  de- 

*  In  generalizing  in  regard  to  these  double  halides  we  must  except  the 
double  cyanides  (page  296). 

t  Compare  the  formulae  Al  P  C18  and  Al  PO4 . 

|  Excepting  with  excess  of  ammonia  solution,  for  aluminium  hydroxide 
is  but  little  soluble  in  that  reagent.  Precipitation  of  aluminium  hydroxide 
from  aluminium  salts  by  means  of  an  excess  of  ammonia  water  is  complete,  if 
the  precipitate  and  liquid  are  boiled  until  the  odor  of  ammonia  disappears. 


ALUMINATES;    SPIKELLS.  339 

rived  from  the  normal  compound  by  loss  of  water  ;  the  most  impor- 
tant of  these  is  the  meta-hydroxide,  Al  02  H,*  which  is  found  as  the 
mineral  diaspor.  The  salts  of  this  substance,  formed  by  replacing 
the  hydrogen  with  a  metal,  are  the  types  of  the  important  group  of 
minerals  known  as  the  spinells.  The  spinells  are  a  group  of  isomor- 
phous  compounds,  each  one  of  which  is  derived  from  a  hydroxide, 

X  -j  QTT   in  which  hydroxide  X  can  be  either  trivalent  aluminium, 


iron,  or  chromium  ;  the  hydrogen  of  this  hydroxide  is  replaced  by  a 
divalent  metal,  M",  so  that  the  general  formula  for  these  minerals 
would  be  M"  (X  02)2;  M"  is  either  divalent  iron,  magnesium,  man- 
ganese, or  zinc.  The  typical  spinell  is  the  aluminate  of  magnesium, 
Mg  (A102)2.  When  meta-aluminium  hydroxide  is  dissolved  in 
caustic  alkalies,  the  solution,  unless  it  is  a  very  concentrated  one 
containing  an  excess  of  the  solvent,  contains  the  meta-aluminate 
of  the  particular  alkali  metal  used  :  — 

Al  02  H  +  MOH  =  Al  02  M  +  H2  0. 

The  aluminates  of  the  alkali  metals  can  be  obtained  in  a  crys- 
tallized state  by  evaporating  a  solution  of  aluminium  hydroxide  in 
an  excess  of  concentrated  caustic  alkali. 

In  addition  to  the  ortho-  and  meta-hydroxides,  another  hydroxide 
of  aluminium,  A12  05  H4  ,  is  frequently  met  with  as  the  mineral 
beauxite  ;  this  substance  is  formed  by  the  separation  of  water 
between  two  formula  weights  of  ortho-aluminium  hydroxide  (see 
page  307  ). 

The  normal  hydroxide  of  aluminium,  when  freshly  precipitated, 
is  a  gelatinous  substance  which  readily  dissolves  in  acids  to  form 
the  salts  of  aluminium  ;  among  these,  perhaps,  the  most  important 
are  the  sulphates. 

ALUMINIUM  SULPHATE.  —  A12  (SO4)3  +  8  H2  O;  formed  by  dissolving  alu- 
minium oxide  or  hydroxide  in  sulphuric  acid  and  evaporating  to  dry- 
ness.  When  heated  it  loses  its  water  of  crystallization  at  100°,  and  at 
red  heat  gives  off  sulphur  trioxide,  leaving  aluminium  oxide  behind  : 
A12  (SO4)3  =  A12  O3  +  3  SO3  . 

ALUMINIUM  AND  POTASSIUM  SULPHATE.  —  A12  (SO4)3  ,  K2  SO4  +  24  H2  O. 
This  substance  is  commonly  known  as  alum.  The  alums  are  double 
salts  composed  of  one  formula  weight  of  the  sulphate  of  a  monovalent 
alkali  metal  or  of  ammonium,  combined  with  the  sulphate  of  a  trivalent 

*  Corresponding  in  formula  to  metaboric  acid,  BO2  H. 


340  ALUMINIUM;    SALTS    OF. 

metal,  the  general  formula  being  M'"2  (SO4)3 ,  M2  SO4  +  24  H2  O. 
M'"  can  be  either  aluminium,  iron,  chromium,  or  one  of  the  rarer 
metals  belonging  to  the  boron  family;  M  can  be  any  one  of  the  alkali 
metals,  or  ammonium.  The  alums  are  all  isomorphous  and  crystallize 
in  octahedra  belonging  to  the  regular  system.  They  can  be  formed 
by  evaporating  to  dryness  a  mixture  of  the  solutions  of  the  sulphates 
of  any  one  of  the  trivalent  metals  mentioned  and  of  one  of  the  alkalies. 

A  number  of  basic  sulphates  of  aluminium  are  known ;  some  of 
these  are  found  in  the  form  of  mineral  deposits.  Several  neutral 
and  basic  phosphates  of  aluminium  occur  in  nature ;  perhaps  the 
most  important  of  these  is  wavellite,  2  Al  P04 ,  Al  (OH)8  +  9  H2  0. 
When  sodium  phosphate  is  added  to  the  neutral  solution  of  an 
aluminium  salt,  the  tertiary  phosphate  of  aluminium,  A1P04,  is 
produced  in  the  ibrm  of  an  insoluble  precipitate ;  this  is  changed  to 
a  basic  phosphate  by  boiling  with  ammonia  water. 

As  has  been  mentioned,  aluminium  is  a  constituent  of  a  very 
large  number  of  silicates,  some  of  which  have  already  been  dis- 
cussed. Of  these  the  most  important  is  undoubtedly  the  hydrated 
tertiary  orthosilicate  which  approximately  has  the  composition 
expressed  by  the  formula  H2  A12  (Si04)2  -j-  H2  0  and  which  is 
known  as  clay  or  kaolin.  This  substance  is  the  result  of  the  dis- 
integration of  feldspar,  or  of  rocks  which  contain  a  large  proportion 
of  that  mineral  (as  some  granites  do)  ;  owing  to  the  destructive 
action  of  the  weather,  the  feldspar  decomposes  into  aluminium  sili- 
cate, silicon  dioxide,  and  the  silicates  of  the  alkali  metals ;  the 
latter  are  washed  away,  brought  into  the  soil,  and  there,  after  in- 
teracting with  other  chemical  constituents  with  which  they  come  in 
contact,  they  are  in  a  proper  condition  to  be  absorbed  by  plants. 
The  kaolin  which  remains  on  the  spot  where  disintegration  occurs, 
owing  to  the  formation  of  silicon  dioxide  during  the  process  of  de- 
struction, is  necessarily  mixed  with  that  substance  and  not  infre- 
quently contains  mica ;  some  clays,  however,  are  washed  to  some 
distance  from  their  place  of  formation,  and  these  may  have  taken 
up  the  most  varied  impurities,  such  as  the  carbonates  of  calcium 
and  magnesium,  the  oxides  of  iron  and  of  manganese,  quartz  sand, 
and  other  materials. 

An  impure  clay,  not  infrequently  colored  by  ferric  oxide,  is 
used  in  the  manufacture  of  bricks ;  it  is  pressed  into  moulds  and 
baked  in  a  kiln  until  it  becomes  hard.  Pure  kaolin  is  white  in 
color  and  is  used  in  the  manufacture  of  porcelain ;  when  it  is  moist 


CLAY  ;   PORCELAIN.  341 

it  forms  a  very  plastic  mass.  Kaolin  which  is  entirely  free  from 
iron  is  alone  useful  in  the  manufacture  of  porcelain  because,  as  the 
mass  is  heated  to  the  point  where  it  softens  and  becomes  glassy, 
any  foreign  substances  would  make  themselves  apparent  by  their 
color.  The  clay  is  purified  by  being  agitated  with  water ;  when  the 
coarser  portion  separates  at  the  bottom,  the  finer  parts  are  moulded 
into  forms,  dried,  and  finally  heated  to  a  red  heat.  This  latter 
treatment  makes  the  article  undergoing  the  process  of  manufacture 
strong,  but  leaves  it  porous;  in  order  to  finish  the  same  it  is  covered 
with  a  mixture  of  silicon  dioxide,  aluminium  oxide,  and  sodium  car- 
bonate, which  ingredients  form  an  easily  fusible  glass,  and  it  is  then 
heated  to  a  temperature  at  which  the  clay  begins  to  soften  and  at  which 
the  glazing  has  been  converted  into  a  coating  of  transparent  glass. 
Fayence  is  made  of  clay  of  somewhat  coarser  structure  than  that 
used  in  the  manufacture  of  porcelain ;  the  thickness  of  the  dishes 
is  greater,  and  the  ware  is  not  heated  to  a  temperature  high  enough 
to  convert  it  into  the  glass-like  mass  which  forms  porcelain.  Fay- 
ence is  covered  with  a  glaze  which  is  much  like  that  given  above, 
with  the  difference  that  some  lead  oxide  is  added.  Common  stone 
ware  is  made  from  clay  which  is  even  more  impure  than  that  from 
which  fayence  is  prepared ;  the  glazing  is  either  put  on  by  covering 
the  ware  with  the  mixture  to  be  used  before  burning,  or  it  is  made 
by  throwing  common  salt  on  the  utensils  while  they  are  being 
heated  in  the  furnace  ;  the  salt  evaporates,  comes  in  contact  with  the 
surface  of  the  materials  used,  and  covers  it  with  a  fusible  soda  glass. 
But  one  sulphide  of  aluminium,  with  a  formula  A12  S3 ,  corre- 
sponding to  the  oxide,  is  known.  This  compound  can  be  produced 
by  heating  a  mixture  of  powdered  aluminium  and  sulphur,  but,  like 
many  sulphides  of  the  not-metals,  it  is  readily  decomposed  by  water, 
yielding  the  hydroxide  of  aluminium  and  hydrogen  sulphide :  — 

A12  S3  +  6  H2  0  =  2  Al  (OH  )3  -f  3  H2  S  ; 

from  this  it  follows  that  when  an  alkaline  sulphide  is  added  to  a 
solution  of  an  aluminium  salt,  aluminium  hydroxide  is  precipitated ; 
so,  for  instance,  the  following  reaction  takes  place  between  alumin- 
ium sulphate  and  ammonium  sulphide  :  — 

3  (NH4),  S  +  A12  ( S04  )3  -f  6  H2  0  =  2  Al  (OH  )3  +  3  (NH4)2  S04 

+  3  H2  S. 
A  similar   reaction  takes  place  when  a  soluble  carbonate  is  added  to 


342  ALUMINIUM   SULPHIDE. 

the  solution  of  an  aluminium  salt,  for,  owing  to  the  extreme  insta- 
bility of  aluminium  carbonate,  the  hydroxide  and  not  the  carbonate 
is  precipitated  :  — 

3  (NH4)2003  +  A12(S04)3  +  6  H20  =  2  Al  (OH)3  +  3  (NH4)2 
S04  +  3  H2  0  +  3  C02 . 

All  of  the  reactions  which  have  just  been  mentioned  illustrate 
the  weakly  basic  character  of  aluminium  •  oxide  and  hydroxide  and 
show  the  close  relationship  existing  between  aluminium  and  the  not- 
metals. 


GALLIUM. 


CHAPTER   XLVIII. 

GALLIUM,    INDIUM,  AND    THALLIUM. 

Gallium;  symbol,  Ga ;  atomic  weight,  69  ; 
Indium;  symbol,  In ;  atomic  weight,  113.7  ; 
Thallium;  symbol,  Tl ;   atomic  weight,  204.18. 

GALLIUM,  indium,  and  thallium  are  very  sparingly  represented  in 
nature ;  they  are  of  scarcely  any  commercial  importance,  so  that  the 
interest  in  them  is  purely  theoretical  in  its  character,  and  is  taken 
because  they  complete  the  family  of  elements  of  which  boron  and 
aluminium  are  the  chief  representatives. 

Gallium  was  discovered  by  Lecocq  de  Boisbaudran  in  zinc- 
blende.*  The  metal  is  hard,  brittle,  and  crystalline  in  its  structure ; 
it  is  scarcely  malleable  or  ductile ;  it  melts  at  30°. 15  and  is  not 
volatile  even  at  a  high  temperature ;  its  specific  gravity  is  5.95.  The 
metal  scarcely  oxidizes  when  exposed  to  the  air,  and  it  is  readily 
obtained  by  electrolysis  of  a  solution  of  the  oxide  in  alkalies ;  it  de- 
composes steam  and  liberates  hydrogen  (see  page  30)  ;  like  alumin- 
ium, it  is  soluble  in  hot,  caustic  alkalies.  The  chief  characteristics 
of  the  compounds  of  gallium  are  given  in  the  following  table :  — 

OXIDES,  Ga  O,  Ga2  O3 .  The  former  is  the  least  stable  of  the  oxides;  it  is 
basic;  the  latter  is  white,  infusible,  reduced  to  the  metal  at  white 
heat  by  a  current  of  hydrogen.  It  is  both  basic  and  acidic,  but  dis- 
solves only  in  the  most  concentrated  caustic  alkalies. 

HYDROXIDE,  Ga  (OH  )3  ,  formed  by  precipitating  the  solutions  of  soluble 
gallium  salts  with  ammonia  water;  it  is  somewhat  soluble  in  an  excess 
of  the  reagent.  When  heated,  the  hydroxide  readily  loses  water  and 
forms  the  oxide. 

CHLORIDES,  Ga  C12  ,  Ga  C13 .  The  former  is  a  solid  which  melts  at  164° 
and  boils  at  535°;  its  vapors  have  a  specific  gravity  which  corresponds 
to  a  molecular  weight  represented  by  the  formula  Ga  C12 .  The  latter 
was  formerly  supposed  to  have  the  formula  Ga2  C16  (see  page  336),  but 
recent  determinations  of  the  specific  gravity  of  the  vapor  of  gallium 
trichloride  show  that  body  to  have  a  molecular  weight  corresponding 

*  Zinc  sulphide,  Zn  S. 


344  INDIUM. 

to  the  formula  Ga  C13  *  at  440°.  The  trichloride  melts  at  75°. 5  and 
boils  at  215°  to  220°.  It  dissolves  in  water,  but  when  the  solution  is 
evaporated  it  is,  in  part,  changed  into  the  basic  chloride,  just  as  is  the 
case  with  the  trichloride  of  antimony  (page  253). 

THE  SALTS  OF  GALLIUM  are  produced  by  dissolving  the  hydroxide  in  the 
various  acids;  the  sulphate,  when  evaporated  with  the  sulphates  of 
the  alkali  metals  or  of  ammonium,  forms  alums.  (See  page  339.) 

The  discovery  of  gallium  in  1875  was  of  especial  interest  because, 
in  the  periodic  system  of  the  elements  as  arranged  by  Mendelejeff 
a  few  years  previous  to  that  time,  an  element  belonging  to  the 
family  of  which  aluminium  is  a  representative  was  found  to  be 
missing.t  An  element,  the  chemical  and  physical  properties  of 
which  should  lie  between  those  of  aluminium  and  indium,  and 
which  would  have  an  atomic  weight  of  approximately  69,  was 
therefore  predicted  by  Mendelejeff,  and  this  prediction  was  subse- 
quently brilliantly  verified  by  Lecocq  de  Boisbaudran. 

The  next  element  of  this  family  is  indium.  Like  gallium,  it 
occurs  in  some  specimens  of  zinc-blende.  The  element  was  dis- 
covered in  1863  by  fceich  and  Eichter.  It  is  white,  with  a  metallic 
lustre  lying  between  that  of  platinum  and  silver ;  it  is  softer  than 
lead,  and  is  very  malleable  and  ductile.  Its  specific  gravity  is  7.4  ; 
it  melts  at  176°,  and  is  somewhat  volatile  at  red  heat ;  when 
heated  to  redness  in  the  air,  it  burns  to  form  In2  03 .  The  chief 
properties  of  its  compounds  are  given  in  the  following  table :  — 

OXIDES,  In  O  and  In.2  Os .  The  former  is  made  by  reducing  the  trioxide 
in  a  current  of  hydrogen ;  it  burns  in  the  air  to  form  In2  Oa .  The 
trioxide,  In2  O8 ,  is  the  most  stable  oxide  of  indium,  and  is  the  one 
corresponding  to  the  typical  oxide  of  the  family ;  it  is  produced  when 
indium  is  burned  in  the  air,  or  when  the  hydroxide  In  (OH  )3  is  heated ; 
this  latter  substance  is  obtained  by  precipitation  from  solutions  of  in- 
dium salts  by  means  of  ammonia  water.  The  oxide  is  easily  reduced 
to  the  metal  by  heating  the  same  in  a  current  of  hydrogen,  the  monox- 
ide appearing  as  an  intermediary  stage  in  this  reduction.  The  trioxide 
and  the  corresponding  hydroxide  are  mainly  basic  in  their  character; 
they  dissolve  in  acids  to  form  the  salts  of  indium ;  they  are,  however, 
also  weakly  acidic,  for  they  are  dissolved  by  the  hydroxides  of  potas- 
sium or  sodium. 

CHLORIDES,  In  C12 ,  In  C13 .     The  latter  is  formed  by  the  action  of  chlo- 

*  Nilssen  and  Petterson;  Comptes  Rendus;  107,  572. 
t  A  similar  gap  was  found  to  exist  in  the  carbon  family;  this  was  subse- 
quently filled  by  the  discovery  of  germanium  (see  page  309). 


THALLIUM.  345 

rine  on  indium ;  it  sublimes  at  440°  without  melting;  its  vapor  den- 
sity corresponds  to  a  molecule  of  the  formula  In  Cls ;  it  dissolves  in 
water  without  change,  but,  on  heating  the  solution,  hydrochloric  acid 
passes  off  and  a  basic  chloride  is  formed.  Indium  chloride  readily 
unites  with  the  chlorides  of  the  alkali  metals  to  produce  double  salts 
corresponding  in  formula  to  those  of  aluminium  (see  page  337). 

SULPHIDE,  In2  S3 ,  is  formed  by  direct  union  of  indium  and  sulphur  at  red 
heat.  A  sulphohydrate  of  indium,  In(SH)g,  is  precipitated  from 
neutral  or  weakly  acid  solutions  of  indium  salts  b$  hydrogen  sul- 
phide; in  this  respect  the  character  of  indium  approaches  that  of  the 
most  pronounced  metals  of  the  preceding  (carbon)  family. 

THE  SULPHATE  OF  INDIUM,  when  evaporated  with  the  sulphate  of  an 
alkali  metal  or  of  ammonium,  produces  an  alum  (see  page 


Thallium,  the  element  having  the  highest  atomic  weight  in  this 
family,  was  discovered  by  Crookes  in  1861,  that  investigator  find- 
ing it  in  the  residues  covering  the  floors  of  the  channels  of  certain 
sulphuric  acid  works ;  since  that  time  it  has  been  discovered  in 
zinc-blende,  iron  pyrites,  and  copper  pyrites.  The  metal  is  white, 
of  crystalline  structure,  and  greatly  resembles  tin ;  it  is  malleable 
and  ductile,  has  a  specific  gravity  of  11.9,  melts  at  290°,  and  boils 
at  a  white  heat.  Thallium  is  readily  oxidized  in  the  air,  and  dis- 
solves in  sulphuric  or  nitric  acid  without  much  difficulty.  The 
high  atomic  weight  of  thallium  is  unfavorable  to  the  expression  of 
a  very  pronounced  chemical  character,  so  that,  as  is  the  case  with 
lead,  it  appears  with  a  number  of  oxides  which  each,  individually, 
resemble  a  different  group  of  elements ;  for  instance,  the  monoxide, 
T12  0,  is,  in  its  chemical  behavior,  very  much  like  the  oxides  of  the 
monovalent  elements  of  the  alkali  family,  while  the  trioxide,  T12  03 , 
falls  into  line  with  the  similar  compounds  of  the  aluminium  group. 
The  characteristics  of  the  most  important  thallium  compounds  are 
given  below :  — 

OXIDES,  T12  O,  T12  O3 ,  Tl  O2 .  The  first,  thallium  monoxide  or  thallous 
oxide,  is  formed  by  the  slow  oxidation  of  the  metal  in  the  air;  it  is  a 
brownish-black  powder,  which  is  soluble  in  water,  forming  the  hydrox- 
ide T1OH;  this  remarkable  solubility  shows  the  resemblance  between 
this  oxide  and  the  oxides  of  the  alkali  metals,  for  the  latter  are  like- 
wise soluble  in  water;  a  solution  of  thallous  hydroxide  has  a  strongly 
alkaline  reaction;  it  neutralizes  acids  to  form  salts,  which,  for  the 
most  part,  are  soluble  in  water  [resemblance  to  the  salts  of  the  alkalies 
(see  latter)].  Thallium  trioxide  (thallic  oxide),  formed  by  heating 
thallium  to  a  red  heat  in  oxygen,  is  insoluble  in  water;  the  hydroxide, 
Tl  O.2  H,  is  formed  by  precipitating  from  a  solution  of  a  thallium  salt 


346  THALLIUM. 

by  means  of  ammonia  water;  neither  the  oxide  nor  hydroxide  has 
acidic  properties;  both  are  oxidizing  agents,  having  a  great  tendency 
to  change  into  thallous  oxide.  A  higher  oxide  of  thallium,  Tl  0% ,  so- 
called  thallic  acid,  is  also  said  to  exist. 

CHLORIDES,  Tl  Cl,  Tl  C13 .  The  first,  thallous  chloride,  is  insoluble  in 
water,  is  precipitated  from  solutions  of  thallous  salts  by  hydrochloric 
acid,  and  very  much  resembles  the  chloride  of  silver  in  appearance. 
(A  larger  work  must  be  consulted  for  a  description  of  the  thallous 
salts.)  'The  trichloride,  T1C18,  formed  by  treating  thallium  with 
chlorine,  is  decomposed  into  thallous  chloride  and  chlorine  when 
heated. 

SULPHIDES.     Two  sulphides  of  thallium,  T12  S  and  T12  S3 ,  are  known. 


ATOMIC   WEIGHTS.  347 


CHAPTER   XLIX. 

THE    DETERMINATION    OF    ATOMIC    WEIGHTS.      DULONG   AND 
PETIT'S   LAW.      THE    LAW    OF    ISOMORPHISM* 

THE  investigations  into  the  gravimetric  composition  of  chemical 
compounds,  which  were  undertaken  at  the  beginning  of  the  century 
and  which  finally  developed  the  laws  of  definite  and  multiple  pro- 
portions, succeeded  not  only  in  establishing  these  purely  empirical 
laws,  but,  as  the  spirit  of  inquiry  in  man  leads  him  to  seek  a  cause 
behind  every  regularly  recurring  phenomenon  or  law,  naturally  an 
explanation  for  the  laws  of  definite  and  multiple  proportions  was 
looked  for,  and,  as  we  have  seen,  found  in  the  atomic  hypothesis. 
In  spite  of  the  subsequent  almost  universal  acceptance  of  these 
laws,  some  chemists  continued  to  doubt  their  exactness ;  indeed,  the 
gravimetric  determinations  of  Dalton's  time  were  too  unsatisfac- 
tory and  varying  to  inspire  much  confidence ;  when,  at  a  later  date, 
Berzelius  subjected  the  work,  which  had  been  done  with  the  purpose 
of  establishing  the  laws  relating  to  the  definite  composition  of 
matter,  to  a  more  exact  revision,  and  so  became  a  firm  believer  in 
their  existence,  no  adequate  reason  for  their  non-acceptance  by 
other  chemists  could  be  advanced ;  but,  when  the  subsequent  discov- 
ery of  an  error  in  Berzelius's  determination  of  the  relative  weight 
with  which  carbon  enters  into  combination  with  other  elements, 
shook  confidence  in  all  of  the  established  rules,  the  atomic  theory 
was  left  in  a.  most  unsatisfactory  condition,  and  it  was  not  until 
1860  that  the  painstaking  and  accurate  work  of  Stas  succeeded  in 
showing  that  the  laws  of  definite  and  multiple  proportions  are  not 
merely  approximate,  but  are,  in  reality,  mathematically  exact.  How- 
ever, the  atomic  hypothesis  had  existed  in  its  present  form  before 
Stas's  time,  and  it  naturally  was  the  endeavor  of  chemists  to  deter- 
mine, not  only  the  mere  fact  that,  for  instance,  a  parts  by  weight 

*  See  also  Lothar  Meyer;  Die  Grundziige  der  Theoretischen  Chemie;  Leip- 
zig, 1890.  (Outlines  of  Theoretical  Chemistry,  trans,  by  Bedson  and  Williams; 
Longmans.) 


348          ATOMIC   WEIGHTS   BY  AVOGADKO'S   HYPOTHESIS. 

of  chlorine  always  unite  with  b  parts  by  weight  of  silver  to  form 
the  chloride  of  silver,  they  also  endeavored,  as  we  have  seen  in  the 
preceding  portions  of  the  work,  by  calling  to  their  aid  various  hy- 
potheses and  theories  of  greater  or  less  plausibility,  to  fix  exactly 
the  relative  weights  of  the  atoms  of  silver  and  chlorine ;  these 
atomic  weights  must  necessarily  bear  such  a  relationship  to  each 
other  that,  in  uniting  to  form  silver  chloride,  they  would  always 
produce  that  substance  with  the  proportion  of  a  parts  by  weight  of 
chlorine  to  b  of  silver.  It  is  evident  that  we  can  only  determine 
the  atomic  weights  from  the  stoi'chiometric  quantities,  provided  we 
have  some  means  of  knowing  the  number  of  atoms  united  to  form 
the  molecules.  This,  however,  it  is  not  possible  to  do  by  direct 
observation,  so  that,  in  selecting  atomic  weights,  we  must  resort  to 
more  or  less  probable  hypotheses.  Gravimetric  determinations 
alone,  therefore,  can  do  no  more  than  give  the  relative  parts  by 
weight  with  which  two  or  more  substances  unite  to  form  a  chem- 
ical compound;  thus,  in  studying  the  composition  of  water  by 
weight,  we  could  but  determine  that  eight  parts  by  weight  of  oxy- 
gen unite  with  one  part  of  hydrogen ;  it  was  only  after  chemists 
combined  the  lesson  taught  by  the  gravimetric  composition  of  water 
with  the  phenomena  attendant  on  its  formation  and  decomposition, 
i.e.,  with  the  facts  that  two  volumes  of  hydrogen  always  unite  with 
one  volume  of  oxygen  to  form  water,  and  that  water,  when  decom- 
posed, always  yields  two  volumes  of  hydrogen  and  one  of  oxygen, 
and  after  these  experimental  facts  were  explained  by  Avogadro's 
hypothesis  (page  70  and  sub.),  that  the  conclusion  was  definitely 
reached  that  each  molecule  of  water,  is,  in  reality,  composed  of  two 
;atoms  of  hydrogen  united  to  one  of  oxygen,  by  reason  of  which 
conclusion  the  atomic  weight  of  oxygen  was  fixed  at  16  and  not  at  8. 
As  has  been  repeatedly  mentioned,  the  determinations  of  the 
specific  gravities  of  gases,  if  we  accept  Avogadro's  hypothesis,  give 
us  the  magnitudes  of  their  molecular  weights,  and,  when  these  are 
fixed,  provided,  in  each  case  the  accurate  sto'ichiometric  *  compo- 
sition of  the  substance  in  question  is  known, t  we  can  determine  the 

*  The  relative  proportions  by  weight  in  which  substances  unite  to  form 
chemical  compounds  are  called  the  stoichiometric  quantities.  (See  page  6.) 

t  It  is  evident  that  the  determination  of  the  vapor  density  of  a  substance 
is  of  no  value  unless  the  stoichiometric  composition  is  known.  For  instance, 
it  is  of  no  influence  on  the  determination  of  the  atomic  weight  of  nitrogen  to 


MAXIMUM   ATOMIC    WEIGHTS. 


349 


maximum  atomic  weights  of  the  elements  entering  into  the  struc- 
ture of  the  molecules  of  the  various  gases.  (See  page  73.)  The 
following  table,  which  illustrates  the  method  by  which  maximum 
atomic  weights  are  determined  from  a  comparison  of  the  vapor 
densities  of  gases,  will  serve  to  more  clearly  fix  these  facts  in  the 
mind  of  the  pupil :  — 


NAME  OF  GAS. 

d. 

d  X  28.8. 

M. 

ANALYSIS  BY  WEIGHT. 

Nitric  oxide. 

1.039 

30. 

30.03 

14.03   nitrogen'      +    16       oxygen. 

Nitrogen  dioxide. 

1.58 

45.5 

46.03 

14.03          "                +    32           " 

Phosphorus  trichloride. 

4.88 

140.9 

137.35 

31.       phosphorus  +  106.35  chlorine. 

Phosphorus  tri-iodide. 

14.46 

417.1 

411.55 

31.                 "           +  380.55  iodine. 

Hydrochloric  acid. 

1.247 

36. 

36.458 

1.008  hydrogen      +    35.45  chlorine. 

Hydroiodic  acid. 

4.443 

128. 

127.858 

1.008          "             +  126.85  iodine. 

Water. 

0.623 

17.99 

18.016 

2.016          "             +    16       oxygen. 

Sulphur  dioxide. 

2.247 

64.9 

64.06 

32.06    sulphur         +    32 

In  this  table  d  is  the  specific  gravity,  air  =  1 ;  d  x  28.8  is  the  specific  grav- 
ity, H2  =  2 ;  *  M  is  the  molecular  weight,  found  by  adding  the  figures  given  in 
the  last  column. 

After  a  study  of  the  above  table  we  can  see  that,  were  chlorine, 
for  example,  to  occur  in  but  one  compound,  and  that  one  the  tri- 
chloride of  phosphorus,  evidently  the  atomic  weight  of  the  element 
would  be  placed  at  a  maximum  of  106.35 ;  it  could  be  no  greater, 
for  the  molecular  weight  of  the  chloride  of  phosphorus  is  known. 
However,  if  we  glance  further  down  the  column,  we  discover  ny- 
drochloric  acid  with  a  molecular  weight  of  36.457  ;  each  molecule 
of  this  contains  but  35.45  parts  by  weight  of  chlorine ;  it  follows 
therefore  that  the  weight  of  106.35,  which  is  the  amount  of  chlo- 
rine contained  in  the  chloride  of  phosphorus,  really  represents 
three  atoms  of  chlorine,  provided  the  molecule  of  hydrochloric  acid 
contains  lut  one  of  these.  The  number  35.45  must  therefore  be 
fixed  upon  as  the  atomic  weight  of  chlorine,  and  must  remain  so, 

determine  the  specific  gravity  of  ammonia  and  find  this  to  be  .589,  and  that 
the  molecular  weight  is  17;  we  must  also  know  that  this  17  parts  by  weight  of 
ammonia  contains  14  parts  by  weight  of  nitrogen ;  when,  however,  the  two  facts 
are  combined,  we  can  say  that  the  atomic  weight  of  nitrogen  cannot  be  more 
than  14,  for,  after  we  have  ascertained  the  specific  gravity  of  ammonia,  we 
then  have  proof  that  there  is  a  substance,  the  molecular  weight  of  which  is 
not  greater  than  17,  which  contains  but  14  parts  of  nitrogen. 

*  The  specific  gravity  of  air  (H.2  =  2)  is  28.8,  hence  specific  gravities  taken 
with  air  as  unity  are  converted  to  those  with  H2  =  2  by  multiplying  by  28.8. 


350  DETERMINATION   OF   MOLECULAR   WEIGHTS. 

unless,  at  some  future  time,  we  were  to  discover  a  compound  of 
chlorine,  the  molecular  weight  of  which  is  known  and  which  con- 
tains relatively  less  than  this  quantity.  In  the  latter  event,  a 
molecule  of  hydrochloric  acid  would  necessarily  contain  more  than 
one  atom  of  chlorine.  Similar  considerations  will  help  us  to  select 
the  number  126.85  as  representing  the  atomic  weight  of  iodine, 
while  a  comparison  of  the  figures  in  the  table  given  above  will 
further  show  us  that,  by  means  of  the  determinations  even  of  the 
very  few  substances  mentioned  there,  the  maximum  atomic  weights 
of  nitrogen,  phosphorus,  sulphur,  oxygen,  chlorine,  and  iodine  are 
given.  These  same  methods  of  investigation  have  been  applied  in 
every  case  where  the  study  of  elements  and  compounds  in  the 
gaseous  state  has  been  possible,  so  that  the  maximum  atomic 
weights  of  the  greater  number  of  elements  have  been  ascertained 
with  reasonable  certainty. 

The  determination  of  the  specific  gravities  of  gases,  although  by 
far  the  most  important,  is  not  the  only  method  for  ascertaining 
molecular  weights.*  In  1882  F.  M.  Raoult  demonstrated  that 
aqueous  solutions  of  organic  substances,  provided  they  contain  the 
dissolved  compounds  in  quantities  proportional  to  their  molecular 
vughts,  have  identical  freezing  points,  and,  subsequently,  the  same 
law  was  found  to  hold  good  for  other  substances  as  well,  although 
the  amount  of  depression  differs  for  each  solvent.  If  A  is  the  low- 
ering of  the  freezing  point  of  a  solvent,  brought  about  by  the  solu- 
tion of  n  molecular  weights  of  a  certain  substance  in  g  grams  of  the 
solvent,  then :  — 

1.  A  =  r-, 

9 

where  r  is  a  constant  depending  only  on  the  nature  of  the  solvent. 
When  the  molecular  weight  of  the  substance  is  not  known,  this  can 
be  ascertained  by  experimentally  determining  the  lowering  of  the 
freezing  point  of  the  solvent  brought  about  when  p  grams  of  the 

substance  are  dissolved  in  g  grams  of  that  solvent,  for  then  n=^-, 

M 
where  M  is  the  molecular  weight,  so  that  equation  1  becomes :  - 

2.  A  =  ^,or, 

Mg 

*  See  Ostwald;  Outlines  of  General  Chemistry  (Walker,  page  136).  The 
description  of  the  method  in  that  book  is  here  given. 


DETERMINATION   OF   MOLECULAR   WEIGHTS.  351 

3.  M^ZL. 

*</ 

The  constant,  r,  can  be  determined  once  and  for  all,  for  any 
given  solvent,  by  dissolving  one  or  two  substances  of  known  molec- 
ular weight  therein,  and  observing  the  depression  of  the  freezing 
point  ;  for,  from  equation  1  :  — 


The  objection  to  this  method  of  determining  molecular  weights 
lies  in  the  fact  that  it  is  applicable  only  in  the  limited  number  of 
cases  where  substances  are  soluble  in  a  medium  which  is  capable 
of  being  frozen  at  a  convenient  temperature,  while,  furthermore,  the 
law  has  not  held  good  in  a  number  of  cases  which  have  been 
observed.*  The  application  of  this  method  is,  nevertheless,  very 
valuable  to  confirm  molecular  weights  determined  by  some  other 
means,  and  to  ascertain  the  same  in  many  cases  where  the  determi- 
nation of  the  specific  gravities  of  gases  is  impracticable.  Similar 
methods,  based  upon  the  lowering  of  the  vapor  pressure  of  solutions 
by  reason  of  dissolved  substances,  have  also  been  applied  to  the 
determination  of  molecular  weights  ;  for  their  study  the  pupil  must 
refer  to  some  text-book  more  especially  devoted  to  these  subject.-  Y 

The  atomic  weights  which,  by  the  application  of  Avogadro's 
hypothesis  (  page  70  )  and  the  determination  of  the  vapor  densities 
of  elementary  and  compound  substances,  have  been  selected  as  those 
which  really  represent  the  relative  weights  of  the  individual  atoms, 
would  be  much  more  worthy  of  confidence  if  other  physical  meas- 
urements could,  when  correctly  interpreted,  indicate  that  this 
selection  had  been  properly  made.  Such  an  aid  is  found  in  the 
application  of  a  law  discovered  by  Dulong  and  Petit  in  the  early 
part  of  this  century.  These  two  investigators  proved  that  the 

*  Salts  are  in  part  decomposed  by  solution  in  the  same  direction  as  they 
are  by  the  electric  current;  so,  for  example,  each  molecule  of  potassium,  so- 
dium, or  ammonium  chloride  is  dissociated  into  a  univalent  not-metallic  atom 
(an  anion),  and  a  univalent  metallic  atom  (or  group  such  as  NH4)  (a  cathion) 
by  great  dilution  ;  the  dissociation  in  the  case  of  these  salts  is  almost  a  complete 
one;  it  is  less  in  the  case  of  sulphates,  etc.  Of  course  the  lowering  of  the 
freezing  point  with  solutions  of  such  substances  [so-called  electrolytes  (see 
page  13  and  28)]  does  not  give  results  indicative  of  the  molecular  weights.  See 
Harry  C.  Jones;  Zeitschrift  fiir  Physikal.  Chemie;  xi.,  535. 

t  Ostwald;  Outlines  of  General  Chemistry;  Walker's  translation. 


352 


LAW    OF   DTJLONG   AND   PETIT. 


greater  the  combining  weight  of  any  given  element  was  found  to 
be,  the  less  was  the  specific  heat  (capacity  for  heat)  of  that  element 
in  the  solid  form,  so  that  the  product  of  the  combining  weight  of  any 
given  element  and  its  specific  heat,  was  found  to  be  very  nearly  equal 
to  the  same  product  for  any  other  element.  This  law  is  susceptible 
of  a  very  simple  physical  explanation.  The  specific  heat  of  a  body 
is  the  quantity  of  heat  necessary  to  increase  the  temperature  of  the 
unit  weight  of  that  body  by  1° ;  it  follows  that  the  product  of  the 
specific  heat  and  the  combining  weight  of  an  element  represents 
the  quantity  of  heat  necessary  to  warm  the  combining  weight 
through  1°,  and  if  we  select  those  numbers  which,  by  using  Avoga- 
dro's  hypothesis,  we  have  decided  upon  as  being  the  atomic  weights, 
then  the  above-mentioned  product  is,  in  the  great  majority  of  cases 
very  nearly  6.4 ;  this  product  can  be  termed  the  "  atomic  heat "  of 
the  elements,  so  that  Dulong  and  Petit' s  law  assumes  the  following 
simple  form :  — 

The  atomic  heats  (capacities  for  heat  of  the  individual  atoms) 
of  all  elements  are  very  nearly  equal* 

This  law  is  true  without  exception  with  all  perfect  (ductile) 
metals  ;  it  holds  good  with  almost  all  metals  which,  like  antimony  or 
tin,  also  have  not-metallic  properties  (which  are  brittle,  but  which 
have  metallic  lustre),  and  also  applies  to  the  greater  number  of  not- 
metals.  A  few  examples  will  serve  as  an  illustration  :  — 


ELEMENT. 

c 

a 

ac 

Lithium. 

.941 

7.02 

6.6 

In  this  table:  — 

Magnesium. 
Chromium. 
Iron. 

.250 
.121 
.114 

24.3 
52.1 

56. 

6.07 
6.3 
6.38 

c    =  specific  heat, 
a    =  atomic  weight, 
ac  =  atomic  heat. 

Cobalt. 

.107 

59. 

6.3 

Nickel. 

.108 

58.7 

6.3 

Bromine. 

.084 

79.95 

6.7 

Gold. 

.032 

197.3 

6.3 

If  we  are,  therefore,  acquainted  with  an  element,  the  maximum 

*  For  a  more  extended  exposition  of  the  results  obtained  by  Dulong  and 
Petit's  law,  see  Lothar  Meyer;  Die  Grundzuge  der  Theoretischen  Chemie; 
Leipzig,  1890.  (Outlines  of  Theoretical  Chemistry,  trans,  by  Bedson  and 
Williams;  Longmans.) 


EXCEPTIONS   TO   LAW   OF   DULONG   AND  PETIT.  353 

atomic  weight  of  which  we  have  never  been  able  to  determine  by 
means  of  the  vapor  densities  of  some  of  its  compounds,  we  can,  as  a 
next  resort,  select  as  its  true  atomic  weight  that  stoichiometric 
quantity  which,  when  multiplied  by  the  specific  heat  of  the  ele- 
ment in  question,  will  give  us  a  number  approximately  equal  to  6.4. 
Of  course,  were  no  method  available  for  fixing  atomic  weights  as 
valuable  as  that  given  us  by  the  determinations  of  the  specific  grav- 
ities of  gases,  this  law  of  Dulong  and  Petit  would  lead  to  no 
definite  results ;  for  it  is  obvious,  when  we  consider  the  above  table, 
that  were  we  to  select  as  atomic  weights  numbers  exactly  one-half 
as  large  as  those  given  above,  the  product  ac  would  still  remain 
constant  and  would  be  approximately  equal  to  3.2 ;  so  that  only  if, 
by  the  use  of  other  physical  and  chemical  means,  we  have  deter- 
mined a  certain  number  of  atomic  weights,  and  have  definitely  de- 
cided that  in  the  case  of  the  elements  in  question,  that  product  must 
be  6.4,  and  not  3.2,  it  then  follows  that  in  all  undetermined  cases 

the  atomic  weights  must   be  nearly  equal  to  the  quotients  of  — — . 

G 

Naturally,  the  numbers  so  obtained  are  not  the  exact  weights  ;  they 
can  only  be  close  enough  to  the  true  numbers  to  show  us  that  the 
atomic  weights  to  be  selected  are  not  one-half  or  twice  or  three 
times  the  stoichiometric  quantities  which  come  closest  to  those 
indicated  by  dividing  6.4  by  the  respective  specific  heats. 

The  few  marked  exceptions  to  Dulong  and  Petit's  law  are  not 
calculated  to  give  any  great  difficulty,  for,  in  cases  where  they 
occur,  the  elements  in  question  form  a  number  of  gasifiable  com- 
pounds, and  consequently  there  is  no  reason  why  the  specific  grav- 
ities of  these  compounds  should  not  definitely  fix  their  maximum 
atomic  weights.  One  example  of  such  an  exception  will  serve  to 
illustrate  this  conclusion.  The  chemical  equivalent  weight  of  car- 
bon is  3,*  its  specific  heat  as  diamond  is.  1.47,  as  graphite,  1.98. 
The  atomic  weight  of  carbon  must  be  some  rational  multiple  of  its 

.  *  By  the  chemical  equivalent  weight  of  an  element  is  meant  that  quantity 
of  the  element  which  will  enter  into  combination  with  one  part  by  weight 
(one  atom)  of  hydrogen,  or  which  will  take  the  place  of  one  atom  of  hydrogen 
in  a  chemical  compound.  The  equivalent  weight  of  carbon  is  3,  because  three 
parts  by  weight  of  carbon  unite  with  one  of  hydrogen.  The  atomic  weight  of 
carbon  is,  however,  4x3  =  12,  for,  by  a  determination  of  the  specific  gravity  of 
the  gas,  it  has  been  decided  that  a  molecule  of  the  simplest  hydrogen  compound 
of  carbon  has  the  formula  CH4 . 


354  MOLECULAR   HEATS    OF   COMPOUNDS. 

equivalent  weight,  so  that,  according  to  the  rule,  we  must  select  a 
number  as  the  atomic  weight  of  carbon  which,  in  addition  to  being 
a  rational  multiple  of  the  equivalent  weight,  will  also,  when  multi- 
plied by  the  specific  heat  of  carbon,  give  a  product  approximately 
equal  to  6.4.  This  product  is  reached  in  the  case  of  diamond  when 
we  take  fourteen  times  the  equivalent  weight,  and  therefore  it  would 
follow  that  the  atomic  weight  of  carbon  is  42 ;  in  the  case  of 
graphite,  however,  we  would  come  to  a  different  result,  for  then 
ac  =  6.5  when  a  =  33.  These  atomic  weights  are  entirely  impos- 
sible, however,  for  we  are  acquainted  with  a  number  of  carbon  com- 
pounds in  which  (the  vapor  density  and  hence  the  molecular  weight 
being  known)  we  relatively  have  only  12  parts  by  weight  of  carbon, 
so  that  there  remains  no  alternative  but  to  believe  that  carbon  pre- 
sents an  exception  to  'Dulong  and  Petit' s  law.  However,  the  spe- 
cific heat  of  carbon  is  much  greater  at  high  temperatures  than  it  is 
at  low  ones,  at  900°  it  is  .459  (ac  would  then  be  equal  to  5.51),  so 
that,  at  a  white  heat,  carbon  would  probably  follow  Dulong  and 
Petit's  law.  These  facts  make  it  apparent  that  specific  heat  deter- 
minations are  of  no  value  in  the  selection  of  atomic  weights  unless 
they  are  made  through  large  intervals  of  temperature,  and  are  then 
found  to  be  constant.  Boron,  silicon,  beryllium,  phosphorus,  and 
sulphur  have,  like  carbon,  very  small  specific  heats.  These  ele- 
ments, which  are  exceptions  to  Dulong  and  Petit's  law,  have  small 
atomic  weights,  and  are  (all  but  beryllium)  not-metals. 

Dulong  and  Petit's  law  is  applicable  in  other  cases  besides  those 
in  which  the  specific  heats  of  solid  elements  are  to  be  considered, 
for,  as  the  capacity  for  heat  of  a  solid  compound  is  very  nearly  equal 
to  the  sum  of  the  capacities  of  its  constituent  parts,  it  is  evident  that, 
in  many  cases,  the  individual  atomic  heats  can  be  obtained  by  a  de- 
termination of  the  specific  heats  of  compounds.  This  method  is 
especially  useful  in  cases  where  the  elements  themselves  are  not  to 
be  obtained  as  solids ;  for  example :  *  - 

Silver;  atomic  weight,  107.9;  specific  heat,  .056;  ac  =    6.04 
Iodine;      "  "        126.8;        "          "       .054;  ac  =    6.8 

The  sum  of  the  atomic  heats  of  silver  and  iodine       =  12.84 

*  Kopp  (Liebig's  Annalen,  Suppl.  3,  290 )  made  a  large  number  of  accurate 
determinations  of  the  specific  heats  of  compound  bodies,  and  to  him  and  Reg- 
nault  is  due  the  development  of  the  proof  of  the  truth  of  law  in  regard  to 
compound  bodies.  The  figures  here  given  are  taken  from  L.  Meyer's  Grund- 
ziige  der  Theoretischen  Chemie,  but  the  original  source  is  to  be  found  in 
Kopp's  researches,  from  which  the  figures  are  taken. 


MOLECULAR   HEATS   OF   COMPOUNDS.  355 

The  observed  specific  heat  of  silver  iodide  is  .0616.  This,  multi- 
plied by  the  sum  of  the  atomic  weights  of  silver  and  iodine,  gives 
14.5,  a  number  which  is  but  little  greater  than  the  one  calculated 
above,  namely,  12.84.  The  same  is  the  case  with  silver  bromide, 
for:- 

Silver  ;     atomic  weight,  107.9;  specific  heat,  .056;  ac  =    6.04 

Bromine;      "  "         79.9;       "  "      .084;  ac  =    6.7 

Calculated  for  Ag  Br  =  12.8 

Silver  bromide;  formula  weight,  187.8;  specific  heat,  .074;    ac  =  13.9 

The  two  examples  given  above  show  us,  therefore,  that  in  the 
case  of  substances  the  formulae  of  which  are  made  up  of  two  atomic 
weights,  the  capacity  for  heat  of  the  formula  weight  of  the  com- 
pound is  very  nearly  equal  to  twice  that  of  either  of  the  individual 
atoms,  and  further  investigation  would  demonstrate  that  compounds 
made  up  of  three  atomic  weights  ( Pb  Br2  and  Pb  I2 )  have  about 
three  times  the  capacity  of  each  individual  atom.  Having  settled, 
by  experimental  evidence,  that  Dulong  and  Petit's  law  holds  good 
for  compounds,  all  of  the  constituent  parts  of  which  can  be  obtained 
in  the  solid  state,  we  can  extend  the  method  so  as  to  indirectly  de- 
termine the  atomic  heats  of  elements  which  are  gases  at  any  tem- 
perature which  permits  of  a  detailed  study  of  their  properties.  For 
example,  we  wish  to  discover  the  atomic  heat  of  chlorine :  — 

Silver  chloride ;  formula  weight,  143.35;  specific  heat,  .0911;  ac  =  13.1 

Silver;  atomic  weight,  107.9  ;         "         "        .056;    ac  =  6.04 

Capacity  for  heat  of  35.45  parts  of  chlorine  =  6.96 

The  difference,  6.9,  obtained  as  the  atomic  heat  of  chlorine  is, 
however,  very  close  to  the  average  of  6.4  which  is  found  by  direct 
observation  in  a  large  number  of  solid  elements ;  so  that  if  we  regard 
the  number  35.45,  as  the  true  atomic  weight  of  chlorine  which  num- 
ber is  selected  by  reason  of  a  large  number  of  specific  gravity  deter- 
minations of  gasifiable  chlorine  compounds,  then  the  atomic  heat  of 
chlorine  very  nearly  coincides  with  the  numbers  obtained  by  direct 
observation  on  a  large  number  of  solid  elements,  and  it  therefore 
follows  that  the  maximum  atomic  weight  of  35.45  is  confirmed  by  an 
application  of  Dulong  and  Petit's  law.  Applications  of  this  method 
with  other  elements  which,  like  chlorine,  have  such  low  melting 
points  that  their  respective  specific  heats  cannot  be  measured  when 
they  are  in  the  solid  state,  have  led  to  like  results,  although  there 
are  a  few  exceptions,  nearly  all  of  which  belong  to  compounds  con- 
taining not-metals  with  small  atomic  weights.* 

*  Compounds  of  hydrogen,  nitrogen,  fluorine,  oxygen. 


356  LAW   OF   ISOMORPHISM. 

Dulong  and  Petit's  law  nelps  us  to  confirm  numbers  already 
decided  upon  as  atomic  weights,  and  when  a  decision  has  been 
reached  in  regard  to  a  large  number  of  elements  by  using  some  other 
method  as  a  guide,  it  can  then  help  us  to  determine  which  numbers, 
representing  multiples  or  simple  fractions  of  the  equivalent  weights, 
are  the  true  atomic  weights  of  elements  which  form  no  gasifiable 
compounds  and  which  are  themselves  not  volatile.  In  no  case  can 
the  law  be  of  assistance  in  determining  molecular  weights. 

A  third  method  which  has  been  of  use  in  determining  the  num- 
bers to  be  selected  as  atomic  weights  has  been  a  study  of  the  laws 
of  isomorphism.  In  1819,  Mitscherlich  discovered  that  certain  ele- 
ments can  replace  others  in  chemical  compounds  without  thereby 
materially  altering  the  crystalline  forms  of  those  compounds. 
These,  as  well  as  the  elements  replacing  each  other  in  them,  are 
said  to  be  isomorphous  (see  page  42).  The  replacing  of  one  ele- 
ment by  another  always  takes  place  in  definite  stoichiometric  quan- 
tities ;  for  example,  in  comparing  the  members  of  the  isomorphous 
group,  represented  by  the  formulae  Na  Cl,  Na  Br,  Na  I,  we  find  that 
35.45  parts  by  weight  of  chlorine  are  always  replaced  by  80  parts 
of  bromine  and  126.85  parts  of  iodine,  and  it  is  very  evident  that, 
provided  the  atomic  weight  of  one  of  these  elements  has  been  decided 
upon  by  some  other  method,  then  the  atomic  weights  of  the  others 
will  be  given  by  the  stoichiometric  quantities  which  can  replace 
that  element  without  altering  the  crystalline  form  of  the  compound. 
This  conclusion  is  correct,  provided  that  in  isomorphous  mixtures 
elements  replace  each  other  atom  for  atom,  and  this  is  the  case  with 
the  vast  majority  of  isomorphous  compounds.*  This  method  for 
determining  atomic  weights  becomes  very  far  reaching  when  we 
consider  that  the  same  element  may  belong  to  two  or  three  isomor- 
phous groups  of  compounds,  so  that,  when  a  decision  as  to  the 
atomic  weights  of  the  elements  in  one  of  these  groups  has  been  ar- 
rived at,  we  can  then  cross  over  to  another,  and,  perhaps,  in  that  one 
discover  an  element  which  also  belongs  to  two  or  three  new  isomor- 

*  Mere  similarity  or  even  identity  of  crystalline  form,  even  when  two 
bodies  which  show  such  similarity  or  identity  are  similarly  constituted,  does 
not  mean  that  the  bodies  in  question  are  isomorphous.  In  order  to  conform 
with  the  laws  of  isomorphism  they  must  be  able  to  replace  each  other  with- 
out thereby  materially  altering  the  form  of  the  crystal.  Such  isomorphons 
compounds  are  the  alums  (page  340),  the  vitriols  (see  magnesium),  the  calcite 
and  arragonite  groups  of  carbonates. 


LAW   OF   ISOMORPHISM.  357 

phous  groups,  and  so  on.  It  follows,  therefore,  that  the  selection  of 
the  atomic  weight  of  one  element  according  to  some  other  method  than 
that  using  the  law  of  isomorphism  may  lead  to  the  determination  of 
the  atomic  weights  of  a  large  number  of  others.  For  example,  the 
atomic  weight  of  zinc,  obtained  by  a  determination  of  the  vapor 
density  of  some  of  the  compounds  of  that  element  which  can  be 
studied  as  gases,*  has  been  placed  at  65.3.  This  proportional  part 
by  weight  of  zinc  when  crystallized  in  the  sulphate  is,  however, 
replaced  by :  — 

58.7  parts  of  nickel ;  56  parts  of  iron ; 
59.        "      "  cobalt ;  55      "      "  manganese,  and 
24.3  parts  of  magnesium. 

Iron,  manganese,  and  magnesium  are  further  isomorphous  with 
40  parts  of  calcium  on  the  one  hand  and  with  27  parts  of  alumin- 
ium and  52.1  parts  of  chromium  on  the  other.  Now,  40  parts  of 
calcium  can  replace  87.6  parts  of  strontium,  137.43  parts  of  barium, 
and  206.95  parts  of  lead,  so  that  this  example  shows  that  the  selec- 
tion of  the  maximum  atomic  weight  of  zinc  can  bring  about  the 
determination  of  the  atomic  weights  of  a  large  number  of  other 
elements,  which  are  connected  with  zinc  by  isomorphous  compounds. 
In  using  the  law  of  isomorphism  one  fact  must,  however,  be  borne 
in  mind.  A  number  of  instances  are  known  in  which  groups  of 
atoms  can  replace  individual  atoms  isomorphously ;  this  is  the  case 
with  ammonium  and  potassium  salts,  for  the  radicle  NH4  takes 
the  place  of  the  atom  K ;  and  if  this  can  be  true  of  ammonium  and 
potassium,  it  is,  without  doubt,  true  of  a  large  number  of  other 
radicles  and  elements;  the  law  can,  therefore,  be  applied  only 
where  no  possible  doubt  can  exist.  All  of  the  methods  which  have 
been  outlined  atbove  have  for  their  object  the  selection  of  those 
multiples  or  fractions  of  the  chemical  equivalent  weights  which 
seem  to  us  to  represent  the  true  atomic  weights.  The  term  "  equiv- 
alent weights  "  was  suggested  at  the  beginning  of  this  century  by 
Wollaston,  because  that  investigator,  as  well  as  a  large  number  of 
others,  was  of  the  opinion  that  no  theories  having  sufficient  plaus- 
ibility would  ever  be  advanced  to  enable  us  correctly  to  determine 
the  true  atomic  weights.  If  we  understand  by  the  expression 
"  equivalent  weights  "  those  quantities  of  the  elements  which 

*  For  example,  zinc  ethyl,  Zn  (C2H5)2. 


358 


EQUIVALENT    WEIGHTS. 


unite  with  one  equivalent,  or  one  part  by  weight  of  hydrogen,  or 
which  can  take  the  place  of  one  part  by  weight  of  hydrogen  in 
compounds,  the  determination  is  comparatively  easy ;  the  true 
atomic  weights  must  then  be  multiples  of  these  equivalent  weights, 
for  one  atom  of  an  element  cannot  unite  with  less  than  one  atom  of 
hydrogen;  this  conclusion  becomes  evident  after  a  study  of  the 
following  table :  — 

One  part  by  weight  of  hydrogen  *  unites  with 


19.      parts  of  fluorine, 
35.45      "     "  chlorine, 
79.95      "     "  bromine, 
126.85      "     "  iodine, 


n. 

8.      parts  of  oxygen, 
16.03      "     "  sulphur, 
39.5        "     "  selenium, 
62.5        "     "  tellurium, 


in. 

4.7  parts  of  nitrogen, 
10.3      "     "  phosphorus, 
25.        "     "  arsenic, 
40.        "     "  antimony, 


and  one  part  by  weight  of  hydrogen  is  replaced  by 


7.02  parts  of  lithium, 
23.05      "     "  sodium, 


39.1 


potassium, 


12.15  parts  of  magnesium, 
20.          "     "  calcium, 
43.8        "     "  strontium, 


9  parts  of  aluminium. 


In  table  I.,  the  atomic  weights  and  the  equivalent  weights  are 
alike;  in  table  II.,  the  atomic  weights  are  twice  the  equivalent 
weights  ;  in  table  III.,  they  are  three  times  the  equivalent  weiglits. 

If  every  element  were  to  combine  with  hydrogen  to  form  a 
well-defined  hydrogen  compound,  or  if  we  could  easily  replace 
hydrogen  in  hydrogen  compounds  by  every  other  element,  the 
determination  of  the  various  equivalent  weights  would  be  compara- 
tively simple.  This  is  not,  by  any  means,  the  case ;  so  that  at 
one  time,  equivalent  weights  were  defined  as  those  weights  which 
would  combine  with  one  equivalent  (eight  parts  by  weight)  of 
oxygen. 

With  this  interpretation  matters  become  much  more  complicated, 
as  the  following  table  will  show  :  — 

8  parts  by  weight  of  oxygen  (one  equivalent)  unite  with:  — 

8.68,  or  17.37,  or  26.05  parts  by  weight  of  chromium, 

or  with :  — 

18.6,  or  21  or  28  parts  by  weight  of  iron,  etc. 

Among  these  weights,  which  is  to  be  selected  as  the  equivalent 
weight  of  iron  or  of  chromium  ?  One  thing  only  appears  certain, 
the  various  equivalent  weights  which  we  have  to  choose  from  in 

*  Although  the  true  atomic  weight  of  hydrogen  is  1.008,  oxygen  being  16, 
the  decimal  can  be  neglected,  and  hydrogen  be  placed  as  =  1. 


ELECTROLYTIC  EQUIVALENTS.  359 

the  case  of  any  given  element,  are  in  simple  ratio  to  each  other,  so 
that  it  follows  that  the  true  atomic  weight  of  any  given  element 
must  be  either  a  simple  multiple  or  a  fraction  of  one  of  its  equiva- 
lent weights  ;  but  which  multiple  or  which  fraction  ?  It  was  from 
such  confusion,  when  many  chemists  had  abandoned  all  hope  of 
ever  selecting  the  true  atomic  weights  of  the  elements,  that  the 
logical  application  of  the  results  obtained  from  Avogadro's  hypo- 
thesis rescued  chemistry,  by  giving  to  it  a  uniform  theoretical  basis 
for  determining  atomic  weights. 

One  method,  which  gave  tolerably  uniform  results  in  the  deter- 
mination of  equivalent  weights,  was  inaugurated  by  Faraday  when 
he  discovered  the  fact  that,  when  an  electric  conductor  of  the  sec- 
ond class  (an  electrolyte)  *  is  made  a  conductor  in  a  circuit,  the 
amount  of  the  electrolyte  which  will  be  decomposed  is  always  pro- 
portional to  the  strength  of  the  current.  If,  then,  we  allow  the 
same  current  to  pass  through  two  electrolytes,  the  constituents 
which  are  separated  from  these  are,  electrically  and  chemically, 
equivalent.  In  order,  therefore,  to  determine  the  equivalent  weight 
of  an  element  by  this  means  we  have  but  to  decompose  some  com- 
pound of  that  element  by  an  electric  current  which,  at  the  same 
time,  traverses  some  conducting  hydrogen  compound  (dilute  sul- 
phuric acid  for  instance),  and  then  to  discover,  experimentally,  what 
part  by  weight  of  the  element  is  separated  simultaneously  with  one 
part  by  weight  of  hydrogen.  This  method  is,  however,  also  sub- 
ject to  error.  For  example,  if  we  decompose  the  chloride  of  iron 
or  of  copper,  we  are  puzzled  in  the  selection  of  the  equivalent  weights 
of  these  elements  because  each  forms  two  chlorides,  which  latter 
compounds  separate  different  quantities  of  metal  in  the  same  time 
and  with  the  same  current ;  the  quantities  of  iron  separated  are,  how- 
ever, in  simple  ratio  to  each  other,  and  the  same  is  true  of  the  quan- 
tities of  copper.  To  give  an  example,  if  we  electrolyze  ferrous 
chloride,  28  parts  of  iron  will  separate  simultaneously  with  one  part 
of  hydrogen ;  but  if  we  electrolyze  ferric  chloride,  18.6  parts  of  iron 
will  be  deposited  during  the  same  time ;  these  quantities  of  iron  are 
to  each  other  as  3  :  2.  Another  objection  to  the  method  of  deter- 
mining equivalent  weights  by  electrolysis  is  that  the  compounds 
of  many  elements  are  non-conductors  of  electricity. 

*  A  substance  which  conducts  electricity  while  being  at  the  same  time 
decomposed,  for  instance,  acidulated  water  (see  page  28). 


360  PRESENT   ATOMIC    WEIGHTS. 

After  the  equivalent  weights  have  been  determined  by  accurate 
methods  of  quantitative  analysis,  and  after  all  hypotheses  and  ex- 
perimental facts  have  been  considered,  we  can  select  the  maximum 
atomic  weights  (which  are  multiples  of  these  equivalent  weights)  be- 
cause of  the  true  understanding  of  the  meaning  of  the  numbers 
obtained  by  the  determination  of  the  specific  gravities  of  gases,  or 
of  molecular  weight  determinations  brought  about  by  other  means 
(and  hence  by  the  methods  founded  upon  Avogadro's  hypothesis). 
These  atomic  weights  have  been  universally  accepted  as  the  basis 
of  our  present  system.  This  solution  seems  to  be  the  correct  one  if 
we  remember  that  only  our  present  atomic  weights,  when  arranged 
in  order,  beginning  with  that  belonging  to  the  element  with  least 
atomic  weight,  and  ending  with  that  with  the  greatest,  form  a  natu- 
ral system,  in  which  elements  with  similar  properties  recur  after 
stated  intervals.  This  arrangement,  known  as  the  periodic  system 
of  the  elements,  will  be  discussed  in  the  next  chapter. 


PERIODIC   SYSTEM.  361 


CHAPTER   L. 

THE   PERIODIC    SYSTEM    OP   THE   ELEMENTS. 

THE  theoretical  basis  for  the  previous  discussions  in  this  book 
has-  been  the  assumption  that  the  properties  of  compounds  are,  in 
reality,  the  properties  of  the  molecules  which,  when  massed  to- 
gether, make  up  those  compounds ;  and  we  have  also  come  to  the 
conclusion  that  the  character  of  any  individual  molecule  depends  on 
the  nature  and  also  upon  the  relative  position  of  the  atoms  which, 
in  uniting,  produce  that  molecule.  If  now,  we  further  wish  to  dis- 
cover whether  any  connection  exists  between  the  characters  of  the 
various  elements,  we  must  first  attempt  to  compare  those  funda- 
mental constants  which  appertain  to  the  atoms  and  which  are  also 
definitely  determinable  or  *to  be  ascertained  with  great  accuracy. 
Such  constants  are  both  the  atomic  weights  and  the  specific  gravi- 
ties of  elementary  bodies. 

In  the  year  1868  both  Mendelejeff  and  Lothar  Meyer*  demon- 
strated that,  when  the  elements  are  arranged  in  the  order  of  their 
increasing  atomic  weights,  each  element  will  differ  in  properties 
from  those  immediately  preceding  and  following  it,  but,  after  certain 
definite  intervals,  elements  will  recur  which  possess  very  similar 
characteristics.  This  interval  they  found,  when  the  elements  are 
arranged  as  indicated,  to  be  after  every  seventh  element  if  the 
first  fourteen  members  alone  are  considered,  but  after  every  seven- 
teenth in  the  remainder  of  the  series.  The  elements  were  therefore 
divided  into  two  "  short  periods  "  of  seven  each,  and  into  five  "  long 
periods  "  of  seventeen  individuals  ;  the  periods  being  so  selected  that 
each  begins  with  an  alkali  metal  and  ends  with  a  halogen ;  for,  by 
this  arrangement,  the  extremities  are  formed  of  sharply  contrasting 
elements.  The  long  periods  are,  however,  in  reality  double,  for  in 
each  one  the  first  seven  elements  resemble  the  last  seven  in  many 

*  Newlands  also  called  attention  to  the  connection  between  atomic  weights 
and  chemical  properties  of  elements. 


362  PERIODIC  SYSTEM;  ARRANGEMENT. 

important  chemical  characteristics,  while  the  three  remaining  ele- 
ments which  are  in  the  middle  *  form  a  separate  group. 

This  arrangement  of  the  elements,  including  the  division  into 
short  and  long  periods,  is  shown  in  the  table  on  the  opposite  page. 

The  first  two  (short)  periods  and  the  third  and  fourth  (long) 
periods  are  complete  with  one  exception,  that  of  an  element  with 
an  atomic  weight  of  about  100  and  corresponding  to  manganese. 
The  remaining  periods  have  only  a  few  known  representatives  in 
each,  and  all  of  the  latter,  with  possibly  one  or  two  exceptions,  may  be 
classed  with  the  very  rare  elements.  As  a  consequence,  all  general- 
izations must  be  made  from  a  study  of  the  complete  periods,  while 
the  conclusions  drawn  from  the  latter  can  only  be  cautiously  ex- 
tended so  as  to  cover  the  ones  in  which  a  number  of  representatives 
are  missing. 

The  short  periods  are  the  types  of  all ;  the  others  are  constructed 
after  those  models  with  this  exception:  While  the  short  periods, 
beginning  as  they  do  with  a  pronounced  metal  and  ending  with  a 
pronounced  not-metal,  will  show  all  gradations  of  character  which 
lie  between  the  two  extremes,  while  traversing  the  five  intervening 
elements ;  the  long  periods,  although  they  begin  and  end  each  with 
an  equally  pronounced  metal  and  not-metal,  nevertheless  require  the 
interposition  of  fifteen  elements  to  effect  the  gradation  which  is  pro- 
duced by  five  in  the  short  periods ;  from  this  it  follows  that  the 
elements  in  the  middle  of  the  long  periods  must  be  further  removed 
in  character  from  those  in  the  middle  of  the  short  periods  than  are 
the  elements  at  either  extremity ;  we  would,  therefore,  expect  ele- 
ments like  chromium,  manganese,  iron,  cobalt,  nickel,  and  copper  to 
have  no  counterparts  in  the  typical  periods.  This  conclusion  is  not 
strictly  correct,  however,  for,  as  has  been  mentioned,  the  long 
periods  are  each  really  composed  of  two  shorter  ones.  The  last 
members  of  the  first  half  of  the  long  periods  therefore  resemble  in 
many  ways  the  last  members  of  the  typical  short  periods,  while  the 
first  elements  of  the  second  half  are  much  like  the  elements  which 
begin  the  short  periods.  Any  given  long  period  begins  with  a  most 
pronounced  metal  like  potassium  or  caesium;  following  this  are 
six  elements,  each  one  of  which  is  less  metallic  in  its  nature  than 
the  one  immediately  preceding  it ;  the  eighth,  ninth,  and  tenth  ele- 

*  Elements  No.  8,  9,  and  10  of  any  long  period. 


PERIODIC    SYSTEM  ;   ARRANGEMENT. 


363 


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364 


PERIODIC    SYSTEM;    ARRANGEMENT. 


ments  then  form  a  transition  from  the  first  to  the  last  seven  of 
the  long  period  ;  they  are  successively  'more  metallic  in  their  nature. 
The  eleventh  element  is  once  more  a  tolerably  pronounced  metal, 
although  it  has  by  no  means  so  metallic  a  nature  as  the  alkali 
metals  which  begin  a  period.  From  the  eleventh  to  the  seventeenth 
element  the  not-metallic  properties  again  become  more  and  more 
pronounced,  until  the  last  elements  of  the  periods  (bromine  or 
iodine)  are  among  the  most  intensely  not-metallic  elements  with 
which  we  are  acquainted.  The  gradation  in  properties  shown  by 
the  periods  can,  perhaps,  be  graphically  represented  as  follows :  — 


Short  period, 
7  elements ;  metallic  properties 

Long  period, 
17  elements;  metallic  properties 


not-metallic 

properties. 


The  elements  which  find  their  places  in  the  vertical  columns 
of  the  periodic  system  form  the  natural  families  or  groups  of  ele- 
ments ;  when  we  compare  the  short  periods  with  the  long  ones  we 
find  that  the  first  two  elements  in  each  of  the  short  periods  belong 
to  the  same  family  as  the  first  two  in  each  of  the  longer  ones,  while 
the  last  five  in  the  former  correspond  to  the  last  five  in  the  latter ; 
as  has  been  pointed  out,  the  intervening  ten  elements  in  the  long 
periods  differ  more  or  less  from  any  which  occur  in  the  typical  pe- 
riods. This  relationship  is  clearly  shown  by  the  following  table  :  — 


Li, 

Be, 

B, 

c, 

N, 

o, 

F, 

Na, 

Mg, 

Al, 

Si, 

P, 

s, 

Cl, 

K, 

Ca, 

Sc, 

Ti, 

V, 

Or, 

Mn, 

Fe, 

Co, 

Ni, 

Cu, 

Zn, 

Ga, 

Ge, 

As, 

Se, 

Br, 

Rb, 

Sr, 

Y, 

Zr, 

Cb, 

Mo, 

—  , 

Itu, 

Rh, 

Pd, 

Ag, 

Cd, 

In, 

Sn, 

Sb, 

Te, 

I, 

Cs, 

Ba, 

La, 

Ce, 

Di, 

—  f 

—  f 

—  t 

—  } 

—  t 

—  , 

—  f 

—  } 

—  t 

—  , 

—  9 

—  , 

Yb, 

—  , 

Ta, 

W, 

—  , 

Os, 

Ir, 

Ft, 

Au, 

Hg, 

Tl, 

Pb, 

Bi, 

—  , 

—  , 

Th, 

—  , 

u, 

—  , 

As  we  pass  from  member  to  member  along  the  complete  series 
of  the  elements  we  encounter,  at  one  time,  a  gradual,  at  another,  an 
abrupt  change  in  the  character  of  the  elements  ;  the  gradual  changes 
between  the  two  extremities  of  any  period,  the  abrupt  changes  as 
we  pass  from  one  period  to  another,  and  these  changes  are  brought 
about,  with  close  resemblance,  in  each  one  of  the  periods,  so  that 
nearly  every  property  of  any  given  element  is  repeated  in  one  or 
more  subsequent  ones;  this  repetition  is  found,  not  only  in  the 
chemical  character  of  the  elements  and  their  compounds,  but  is  also 


ATOMIC   VOLUMES.  365 

apparent,  even  in  a  greater  degree,  in  their  physical  properties,  the 
properties  of  any  individual  element  are  therefore  determined  by  the 
position  of  that  element  in  the  periodic  system. 

One  of  the  easily  determined  constants  which  appertains  to  each 
element  is  its  specific  gravity  in  the  solid  state,  and  that  property 
periodically  increases  and  decreases  as  we  pass  in  regular  order  from 
element  to  element  along  the  entire  series.  This  fact  becomes  very 
apparent  if,  instead  of  comparing  the  specific  gravities  themselves 
(i.e.,  the  quantities  of  matter  contained  in  the  unit  volume),  we 
compare  the  volumes  which  are  occupied  by  weights  of  the  respec- 
tive elements,  which  weights  are  so  taken  that  the  number  of 
grams  correspond  to  the  atomic  weights.  By  this  means  we  can 
arrive  at  the  volumes  occupied  by  the  same  number  of  atoms  in 
each  case,  and  these,  necessarily,  must  bear  the  same  relationship  to 
each  other  as  do  the  volumes  occupied  by  the  individual  atoms. 
These  volumes  can  appropriately  be  termed  "  atomic  volumes,"  and 
they  are  readily  ascertained,  in  the  case  of  each  element,  by  dividing 
the  atomic  weight  by  the  specific  gravity,  so  that :  — 

V  =-. 
c 

In  this  equation  V  represents  the  atomic  volume,  a  the  atomic 
weight,  and  c  the  specific  gravity  of  the  solid  element.  For  in- 
stance, the  atomic  weight  of  lithium  is  7.02,  its  specific  gravity 

7  02 

is  .59,  its  atomic  volume  is  therefore  — —  =  11.9 ;  the  quotient 

.59 

11.9  means  that  7.02  grams  of  lithium  occupy  11.9  cubic  centi- 
metres of  space ;  the  atomic  weight  of  manganese  is  55,  its  specific 
gravity  is  8,  its  atomic  volume  6.9,  therefore  55  grams  of  manga- 
nese take  up  6.9  cubic  centimetres.  The  table  on  following  page 
demonstrates  the  relationship  between  the  atomic  volumes  and  the 
periodic  system. 

In  each  period,  whether  it  be  short  or  long,  the  specific  gravity 
begins  with  a  minimum  (with  the  specifically  light  alkali  metals), 
advances  to  a  maximum  at  the  middle,  and  then  once  more  dimin- 
ishes to  a  minimum  at  the  opposite,  not-metallic  extremity ;  each 
period,  therefore,  represents  a  complete  wave  in  regard  to  specific 
gravities,  the  beginning  being  in  tfre  trough,  the  middle  at  the  crest, 
and  the  end  in  the  succeeding  trough.  The  reverse  is  true  in  regard 
to  atomic  volumes ;  these  begin  with  their  maximum  at  the  alkali 


366 


ATOMIC   VOLUMES. 


metals,  diminish  to  a  minimum  at  the  centre  of  the  periods,  and 
then  once  more  increase  to  a  maximum  at  the  other  extremity ;  the 


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changes  in  the  atomic  volumes  in  each  period,  therefore,  may  be 
compared  to  a  wave  the  crest  of  which  corresponds  to  the  beginning 


ATOMIC    VOLUMES.  367 

alkali  metal,  the  trough  to  the  middle  of  the  period,  and  the  succeed- 
ing crest  to  the  next  following  alkali  metal.  The  atomic  volumes 
of  the  alkali  metals,  however,  increase  rapidly  with  increasing  atomic 
weight,  so  much  so,  indeed,  that  the  atomic  volume  of  sodium  is 
twice  that  of  lithium,  and  the  atomic  volume  of  potassium  is  twice 
that  of  sodium  ;  each  succeeding  wave  which  represents  the  changes 
in  the  atomic  volumes  of  the  elements  forming  one  period  has  a 
greater  amplitude  than  the  one  preceding  it,  and  a  lesser  amplitude 
than  the  one  following.  If,  then,  we  take  the  atomic  volumes  as  our 
guide,  representing  these  as  ordinates  and  the  atomic  weights  as 
abscissae,  we  can  represent  the  periodic  system  in  the  form  of  succes- 
sive wave-like  curves,  the  relative  position  of  any  element  upon 
these  curves  determining  the  properties  of  that  element.  Those 
elements  which  find  their  places  on  a  descending  branch  of  one  of 
these  curves,  and  which  immediately  follow  a  maximum,  and  the 
next  following  elements  down  to  a  minimum  and  even  a  little 
beyond  this  point,  are  difficult  to  fuse,  and  are  not  volatile,  and, 
furthermore,  they  are  less  fusible  the  nearer  they  approach  a  mini- 
mum ;  those  elements  in  the  ascending  curves  are  easily  melted,  and, 
with  few  exceptions,  are  volatile ;  of  the  elements  belonging  to  the 
latter  class,  nitrogen,  oxygen,  and  fluorine  in  the  first  period  are 
gases ;  in  the  second,  chlorine  only  is  a  gas,  phosphorus  and  sulphur, 
however,  melt  at  a  low  temperature  and  are  easily  volatilized  ;  in  the 
first  long  period  the  volatile  elements  begin  with  zinc,  in  the  next 
following  with  silver,  and  in  the  next  (incomplete)  one  with  mercury. 
Lothar  Meyer,  in  consideration  of  these  facts,  has  established  the 
following  rule :  In  that  portion  of  the  series  in  which  the  atomic 
volumes  are  decreasing  with  increasing  atomic  weights,  the  elements 
are  not  volatile  and  are  fusible  with  difficulty ;  on  the  other  hand, 
where  the  atomic  volumes  are  increasing  the  elements  are  easily 
melted  and  are  volatile.  All  of  the  other  properties  of  the  elements 
vary  twice  in  all  of  the  long  periods ;  so,  for  instance,  the  alkali 
metals  at  the  maxima  of  the  curves  as  well  as  the  metals  immedi- 
ately following  are  malleable  and  ductile ;  then,  as  the  minima  are 
approached,  there  follow  brittle,  crystalline  metals ;  these,  in  turn, 
at  the  minima  give  way  to  malleable  and  ductile  ones  ;  and  succeed- 
ing the  latter,  as  the  next  maximum  is  approached,  the  not-metals, 
which  are  neither  malleable  nor  ductile,  find  their  places. 

If  we  designate  as  positive  those  elements  whose  oxides,  in  the 
greater  number  of  cases,  act  as  bases,  and  as  negative  those  ele- 


368 


VALENCE    AND   THE   PERIODIC    SYSTEM. 


ments  whose  oxides  are  anhydrides,*  then  the  periods  all  begin 
with  strongly  positive  alkali  metals ;  next  following,  on  descending 
curves  of  the  atomic  volumes,  are  a  number  of  less  positive  elements  ; 
as  a  minimum  is  approached  the  latter  give  way  to  one  or  two  neg- 
ative (or  at  least  in  greater  part  negative)  individuals ;  succeeding 
these,  at  the  minimum  and  at  the  beginning  of  the  ascent  toward 
the  maximum  are  a  number  of  positive  elements ;  and  finally  the 
curves  are  completed  by  elements  which  are  entirely  negative. 

That  the  chemical  properties  of  the  elements  are  altogether  in 
harmony  with  the  periodic  system  has  been  repeatedly  noted,  as 
certain  sections  of  that  system  have  been  discussed  during  the 
progress  of  this  work ;  and,  as  these  relationships  have  already  been 
explained  at  some  length  at  those  portions  of  the  work  where  the 
various  families  have  been  taken  up,  and  as  they  will  be  further 
noted  in  those  which  are  to  come,  it  seems  scarcely  necessary  at 
this  place  to  do  more  than  to  state  briefly  sonle  few  connections. 

By  determining  the  specific  gravities  of  the  gases  obtained  by 
volatilizing  the  halogen  and  hydrogen  compounds  of  the  elements, 
we  are  able  to  determine  the  respective  valences  of  the  atoms  (see 
page  107).  If  we  compare  the  halogen  and  hydrogen  compounds 
in  the  first  period  with  those  in  the  second,  we  find  the  changes  in 
valence,  as  we  go  from  left  to  right,  to  be  identical  in  both.  The 
beginning  elements  (alkali  metals)  are  invariably  univalent.  Pass- 
ing from  these  to  the  right,  we  find  that  the  valence  is  increased  by 
one  with  each  pair  of  successive  elements  until  a  maximum  of  four 
is  reached  in  connection  with  the  members  of  the  carbon  family ;  it 
then  successively  diminishes  until  it  once  more  reaches  a  minimum 
of  one  (halogen  family).  These  changes  are  made  apparent  by  the 
following  table  :  — 


(H  represents  an  atom  either  of  hydrogen  or  of  chlorine.}  f 

Compounds, 

Li, 
LiK, 

Be, 
BeR2, 

B, 
BR8, 

c, 

CR4, 

N, 
NR3, 

o, 

OR2  , 

F, 
FR. 

Valence, 

1, 

2, 

3, 

4, 

3, 

2, 

1. 

Compounds, 

Na, 
NaR, 

Mg, 
MgRs, 

Al, 
A1R3, 

Si, 
SiR4, 

P, 
PR3, 

s, 

SR2, 

01, 

C1R. 

*  See  page  13. 

t  This  comparison  is  legitimate  only  in  the  first  two  periods,  and  then 
with  the  understanding  that,  when  a  hydrogen  compound  of  an  element  does 
exist,  we  study  that  substance  and  not  the  chloride. 


VALENCE   AND   TH£   PERIODIC    SYSTEM. 


369 


A  different  result  becomes  apparent  in  comparing  the  compounds 
which  the  elements  form  with  oxygen.  The  valence  toward  that 
element  (page  111)  begins  with  one  in  the  family  of  the  univalent 
alkali  metals  and  increases  to  a  maximum  of  seven  in  that  of  the 
halogens.  This  change  becomes  evident  if  we  represent  a  single 
valence  of  one  of  the  atoms  of  oxygen  by  r,  0  =  2r,  and  then  group 
the  elements  as  shown  in  table  :  — 


1 

Li, 

Be, 

B, 

c, 

H, 

o, 

F, 

Oxides, 

Li.20, 

BeO, 

B203, 

CO,, 

N205, 

i 

•> 

Lir, 

Ber2, 

Br3, 

Cr4, 

Nr5, 

~~i 

. 

Na, 

Mg, 

Al, 

Si, 

P, 

s, 

Cl, 

Oxides, 

Na20, 

MgO, 

A1203, 

Si  O,  , 

P205, 

S03, 

(C1207), 

Nar, 

Mgr.2, 

Alr3, 

Sir4, 

Pr5, 

Sr6, 

Clr7. 

Valence, 

1, 

2, 

3, 

4, 

5, 

6, 

7. 

In  making  this  comparison  we  must  remember  that  carbon, 
nitrogen,  phosphorus,  sulphur,  and  chlorine  each  forms  a  number 
of  oxides,  so  that  it  is  only  by  comparing  those  compounds  which 
contain  the  greatest  amount  of  oxygen  that  the  elements  in  a  short 
period  present  a  regular  increase  of  valence  from  one  to  seven. 

In  the  long  periods  there  are  but  few  hydrogen  compounds, 
and  these  belong  only  to  the  last  few  elements  of  the  periods ;  so 
that  if  we  wish  to  compare  the  long  periods  with  the  short  ones,  we 
shall  be  compelled  to  resort  almost  exclusively  to  the  chlorine  com- 
pounds. If  the  halides  of  the  elements  in  the  short  periods  were 
to  correspond  exactly  to  the  hydrogen  compounds,  we  should  be 
able  to  construct  the  following:  table  :  — 


Hydrogen  compounds,  N  H3 , 
PH3, 

Halogen  compounds,  N  C13 , 
PCL, 


OH2, 
SH2, 


FH, 
C1H. 


OC12,     FC1, 
SC12,     C1C1. 


All  of  these  chlorine  compounds  do,  in  reality,  exist ;  but,  in 
addition,  phosphorus  forms  a  pentachloride,  P  C15 ,  and  sulphur  a 
tetrachloride,  S  C14 ,  and  a  monochloride,  S2  C12 ;  and  the  difficulty 
of  a  systematic  comparison  of  the  halides  is  further  enhanced 
by  the  discovery  of  chlorides  having  the  formulae  Mo  C15  and  W  C16 . 
From  these  facts  we  are,  perhaps,  justified  in  drawing  the  conclu- 
sion that,  probably,  in  the  two  short  periods  the  valence  toward 


370 


VALENCE   AND   THE   PERIODIC    SYSTEM. 


chlorine  would  increase  from  one  to  seven,  just  as  it  does  toward 
oxygen,  if  it  were  not  for  the  fact  that  the  not-metallic  elements 
in  those  periods  are  too  negative  to  enable  them  to  retain  any 
great  number  of  chlorine  atoms  in  a  stable  molecule.  This  objec- 
tion does  not  appertain  to  molybdenum  and  tungsten,  and  with 
these,  as  the  above  formulae  show,  the  valence  toward  chlorine  can 
be  five  and  six.  The  dual  nature  of  the  long  periods  is  plainly 
demonstrated  by  a  study  of  the  valence  which  the  elements  belong- 
ing therein  develop  toward  oxygen.  The  initial  alkali  metal  is  al- 
ways univalent ;  advancing  from  this  the  valence  steadily  increases 
with  each  successive  element,  until  a  maximum  of  seven  is  reached ; 
it  then  diminishes  with  the  eighth,  ninth,  and  tenth ;  is  again  equal 
to  one  with  the  eleventh,  and  from  this  to  the  seventeenth  increases 
to  a  maximum  of  seven.  In  order  to  illustrate  these  changes,  we 
will  select  the  first  long  period,  and,  in  order  to  demonstrate  more 
clearly  the  existing  resemblance  between  its  first  and  second  halves, 
will  place  the  one  under  the  other  :  — 


I. 

II. 

III. 

IV. 

V. 

VI. 

VII. 

VI. 

IV. 

III. 

g    K, 

Ca, 

Sc, 

Ti, 

V, 

Cr, 

Mn, 

Fe, 

Co, 

Ni, 

3    KaO, 

CaO, 

Sc203, 

Ti  O2  , 

V205, 

Cr03, 

Mn2O7, 

FeO,, 

Co02, 

Ni203  , 

&    Kr, 

Car2, 

Scr3, 

Tir4, 

VrB, 

Crre, 

Mn  r7  , 

Fer6, 

Cor4, 

Ni  ra  . 

8    Cu, 

Zn, 

Ga, 

Ge, 

As, 

Se, 

Br, 

2j   Cu20, 

ZnO, 

Ga203, 

GeO2, 

As2O8  , 

Se03, 

(Br207), 

0    Cur, 

Znr2  , 

Gar3, 

Ger4, 

Asre  , 

Sere, 

Brr7. 

The  elements  in  the  first  half  of  this  period  (beginning  with 
a  very  intensely  metallic  alkali  metal  and  ending  with  a  metal 
[manganese]),  are  necessarily  much  more  positive  in  their  char- 
acter than  are  those  of  the  second  half,  which  ends  with  a  pro- 
nounced not-metal  (bromine);  nevertheless,  the  two  sections  bear  a 
striking  resemblance  to  each  other.  This  period  is,  therefore, 
formed  of  a  primary  and  secondary  short  period  and  of  three  ele- 
ments, iron,  cobalt,  and  nickel,  which  connect  the  two  halves  and 
form  a  gradual  transition  from  one  to  the  other.  Both  the  primary 
and  secondary  short  periods  resemble  the  typical  short  periods.  This 
relationship  is  made  apparent  in  the  following  table,  which  is  the 
one  in  general  use  :  — 


MENDELEJEFF'S  TABLE. 


371 


o 


DC 


o 

to 


era 


Ci 


OP 

Or 


to 

O 


OP 


o 


p.  Or   -j  .^  OQ 

bO  CO  CO 


to 


00 


tO 

CO    ij 
^    3    fcO 

o  ^  o  >fl 

pro* 

CD 


CO 


OO 


to     . 


03          e 

*      o 


co 


01 


tO 


0 


SO   5^  Cn 

^  Oi    &  fcO 

to  H          °  ^ 


Ci 


01 


0)  O    CD 


O 


CO 


o 


00 


00   l-l 

Or 


s 

Ci   ^ 


*  Now  separated  into  neodymium  (140.5)  and    praseodymium  (143.6). 
Atomic  weight  of  tellurium  is  doubtful  (see  page  104 


372 


TABLE   OF    OXIDES. 


The  elements  which  are  placed  between  the  vertical  lines  consti- 
tute the  natural  families,  those  in  the  horizontal  lines  are  the 
periods  or  series  ;  those  periods  which  have  even  numbers  constitute, 
with  the  exception  of  the  first  short  period  (number  2),  the  first 
sections  of  the  long  periods;  those  with  the  odd  numbers,  with  the 
exception  of  the  second  short  period  (number  3)  form  the  second  sec- 
tions ;  the  series  having  even  numbers,  therefore,  bear  the  closest 
resemblance  to  each  other,  while  those  having  odd  numbers  also  show 
a  great  similarity  of  characteristics  ;  on  the  other  hand,  the  periods 
numbered  with  odd  numbers  bear  a  much  less  marked  resemblance 
to  those  with  even  ones.  These  peculiarities  will  be  demonstrated 
more  at  length  during  the  discussion  of  the  individual  families  of 
metals.  If  a  table  of  the  oxides  of  the  elements  is  constructed, 
while  following  out  the  arrangement  given  above,  the  remarkable 
regularity  displayed  in  the  formation  of  those  compounds  is  made 
apparent.* 


2. 
3. 
4. 
5. 
6. 
7. 
8. 
9. 
10. 
11. 
12. 

i. 

ii. 

III. 

IV. 

V. 

VI. 

VII. 

Li2  0, 
Na20, 
K2    0, 
Cu2  O, 
Rb20, 
Ag20, 
Cs2  0, 
? 

Be202, 
Mg202, 
Ca2  02, 
Zn202, 
Sr2  02, 
Cd2  O2  , 
Ba2  O2  , 

B2    03, 
A1203, 
Sc2  03, 
Ga2  03  , 
Y2   03, 
In2  03  , 
La2  O3  , 

C2    04, 
Si2    O4, 
Ti2  04, 
Ge204, 
Zr2   04, 
Sn2  O4, 
Ce2  O4, 
i 

Pb204, 

No    O5, 

P2"  05, 
Y2  06, 

As205, 
Cb205, 
Sb205, 
Di205, 

Ta205, 
Bi2  05  , 

82     06f 
Cr2  06, 
Se2   06, 
Mo206, 
Te2  06, 

W2  06, 
U2    00, 

C12    07, 
Mn2  O7  , 
Br,  07, 

I2      07, 

i 

Yb,03, 
Tla  03  , 

^ 

> 

5 

Au.20, 

5 

Hg2  O2  , 

5 

! 

In  constructing  this  table  the  formulae  of  the  oxides  in  the 
families  of  beryllium,  carbon,  and  sulphur  have  been  doubled,  so  as 
to  render  the  increase  in  the  valence  toward  oxygen,  as  the  series 
proceed  from  left  to  right,  more  apparent.  Of  course,  a  number  of 
the  elements  form  oxides  with  formulae  differing  from  those  given 
in  the  table ;  only  those  oxides  have  been  selected  for  purposes  of 
comparison,  which  are,  in  any  given  family,  common  to  all  of  the 
members  of  that  family,  or,  in  cases  where  the  oxides  themselves  f 

*  See  Lothar  Meyer;  Grundziige  der  Theoretischen  Chemie;  p.  65. 
t  For  instance,  oxides  of  the  formulae  L>  O7  and  Br2O7  do  not  exist;  but 
the  acids  derived  from  these,  and  the  salts  of  these  acids,  are  known. 


PREDICTION    OF   ELEMENTS. 


373 


are  not  known,  their  existence  has  been  considered  as  theoretically 
possible  because  some  derivatives  of  the  missing  oxides  have  been 
described.  In  any  given  family,  any  one  of  the  oxides  given  on 
the  above  table  in  the  vertical  column  belonging  to  that  family, 
may  be  termed  the  typical  oxide  of  that  group. 

By  a  skilful  combination  of  the  connections  which  have  been 
emphasized  in  the  last  chapter,  Mendelejeff  was  able  to  predict  the 
existence  of  a  number  of  elements,  unknown  at  the  time  of  the  dis- 
covery of  the  periodic  system ;  for  the  purpose  of  illustrating  the 
method  adopted  by  that  investigator,  one  example  will  be  given  here. 

No  element  fitting  into  the  fourth  series,  group  three,  was 
known  at  the  time  when  the  periodic  system  was  discovered ;  yet, 
Avere  such  a  one  to  be  isolated  in  the  future,  it  should,  in  its  proper- 
ties, be  related*  to  aluminium  in  the  same  way  as  calcium  is  to 
magnesium,  or  as  titanium  is  to  silicon.  Its  atomic  weight  should 
be  about  44,  inasmuch  as  it  would  follow  K  (39),  Ca  (40),  and  be 
followed  by  Ti  (48)  and  V  (51).  In  predicting  the  properties  of 
this  element,  Mendelejeff  reasoned  that  it  would  be  as  much 
more  metallic  than  aluminium,  as  calcium  is  than  magnesium,  or 
as  titanium  is  than  silicon.  This  unknown  element  Mendelejeff 
called  ekaboron,  with  a  symbol  Eb ;  and  for  the  purposes  of  com- 
parison, the  predicted  properties  of  ekaboron  and  the  real  proper- 
ties of  scandium,  the  element  which  was  subsequently  discovered, 
are  placed  side  by  side :  *  — 


SCANDIUM. 


Atomic  weight  about  44. 

Oxide,  Eb^  O3  ,  soluble  in  acids,  analogous 
to  AlaO3  ,  but  more  basic;  insoluble  in  alka- 
lies. 

Salts  of  Eb,  colorless,  yield  gelatinous  pre- 
cipitates with  Na  OH,  Na2  Co3  . 

Sulphate,  #&2(S04)3,  will  form  a  double 
salt  with  K2  SO4  ,  which  will  not  be  isomor- 
phous  with  the  alums. 


Atomic  weight,  44. 

Oxide,  Sca  O3  ,  soluble  in  strong  acids,  anal- 
ogous to  AlaO3,  but  decidedly  more  basic; 
insoluble  in  alkalies. 

Salts  of  Sc  are  colorless,  and  yield  gelatin- 
ous precipitates  with  NaOH,  Na2  CO3  . 

Sulphate,  Sc2  (S04)3 ,  forms  a  double  salt 
with  K2  SO4  ,  which  is  not  isomorphous  with 
the  alums. 


It  seems  scarcely  necessary  to  enter  into  a  more  detailed 
description  of  the  periodic  system  at  this  place ;  the  elements 
which  have  already  been  considered  have  been  discussed  in  their 
relation  to  the  natural  groups  of  which  they  are  members,  so  that 


*  See  Pattison  Muir;  Principles  of  Chemistry;  p.  201. 


374  PREDICTION   OF    ELEMENTS. 

their  individual  connections  have  been  sufficiently  pointed  out ; 
those  which  are  to  follow  will  be  described  in  the  order  given  by 
the  periodic  system,  while  attention  will  be  called  to  the  character 
of  the  various  families  at  the  proper  place.  It  does  not  fall 
within  the  scope  of  this  work  to  give  a  detailed  description  of 
each  individual  metal,  as  has  been  done  with  the  not-metals ;  for 
the  ground  is  very  abundantly  covered  by  the  large  number  of 
works  on  qualitative  analysis,  which  discuss  many  of  the  chemical 
reactions  peculiar  to  the  metals  because  the  latter  are  mainly  inter- 
esting from  an  analytical  standpoint ;  and,  moreover,  much  of  the 
chemistry  of  the  salts  of  the  metals  is  merely  a  repetition  of  what 
has  already  been  taken  up. 


NEUTRALIZATION.  375 


CHAPTER   LI. 

NEUTRALIZATION.      DOUBLE    DECOMPOSITION.      DISSOCIATION 
OF    ELECTROLYTES. 

THE  phenomena  attending  the  neutralization  of  an  acid  by  a  base 
are  of  such  importance  that  a  brief  discussion  of  their  nature  is 
necessary.  If,  to  use  potassium  hydroxide  as  an  example,  that  base 
is  brought  in  contact  with  hydrochloric  acid,  the  following  change, 
as  expressed  by  our  chemical  equations,  takes  place :  — 

KOH+  HC1=KC1+H20. 

Of  the  two  systems,  KOH  +  H  Cl  and  K  Cl  +  H2  0,  the  former  is 
in  unstable  equilibrium ;  the  two  systems  correspond  to  two  differ- 
ent quantities  of  energy,  so  that  when  the  former  is  converted  into 
the  latter,  energy  is  conducted  away  in  the  form  of  heat.  The 
sum  of  the  energy  thus  conducted  away,  and  of  that  remaining  in 
the  system  K  Cl  •+  H2  0,  must  be  equal  to  that  originally  contained 
in  KOH  -f  H  Cl.  We  are  also  acquainted  with  chemical  reactions 
in  which  energy  must  be  conducted  to  a  system  in  order  to  change  it 
into  a  second  one ;  the  former  class  of  reactions  are  exothermic, 
the  latter  are  endothermic  (page  12).  In  how  many  portions  the 
energy  may  be  communicated  or  may  pass  off  has  evidently  no 
effect  on  the  final  value.* 

In  the  reaction  cited  above  there  must  be  an  expenditure  of 
energy  sufficient  to  decompose  KOH  into  K  -f  OH  and  H  Cl  into 
H  -f  Cl  before  the  rearrangement  into  K  C1.+  H2  0  can  take  place ; 
but  during  the  complete  reaction  energy  passes  off  in  the  form  of 
heat,  and  the  amount  of  the  latter  is  evidently  independent  of  any 
intermediate  changes.  The  reaction  between  potassium  hydroxide 
and  hydrochloric  acid,  taking  into  account  all  of  the  thermal  values> 
would  be  expressed  as  follows :  — 

KOH  +  H  Cl  =  K,  Cl  [1012  K]  +  H2 , 0  [684  K]  —  (K,  O,  H  [1165  K] 
+  H  Cl  [393  K])  ; 

*  See  Ostwald,  Outlines  of  General  Chemistry,  p.  368. 


376  HEAT    OF   NEUTRALIZATION. 

and  from  this  it  will  be  seen  that  the  sum  of  the  heats  of  forma- 
tion of  a  formula  weight  of  potassium  chloride  and  one  of  water  is 
greater  by  138  K  than  that  of  the  sum  of  the  heats  of  formation  of 
similar  quantities  of  potassium  hydroxide  and  hydrochloric  acid. 
This  may  be  expressed  in  the  following  terms  :  "  The  heat  of  for- 
mation of  K,  Cl  +  H2,  0  is  greater  than  that  of  K,  0,  H  -f  H,  Cl,"  * 
and,  as  we  may  disregard  the  intermediary  changes,  we  may  express 
the  final  result  as  follows  :  — 

KOH  aq  +  HC1  aq  =  K  Claq  +  H2  0  +  aq  +  138  K, 

the  symbol  aq  signifying  that  the  constituents  are  dissolved  in  a 
quantity  of  water  so  large  that  the  addition  of  any  more  of  the 
reagents  will  not  affect  the  thermal  value ;  or,  as  the  formation  of 
water  takes  place  in  all  neutralizations  of  acids  with  bases,  and  as 
its  mixture  with  the  salt  solution  can  produce  no  thermal  effect,  we 
may  write  the  equation  :  — 

KOH  aq  +  H  Cl  aq  =  K  Cl  aq  +  138  K. 

Now,  it  has  been  shown  that  different  acids,  as  well  as  different 
bases,  evolve  different  amounts  of  heat  when  neutralized,  but  the 
difference  between  the  amounts  of  heat  given  off  by  any  two  bases 
when  neutralized  by  a  series  of  acids  or  between  any  two  acids  when 
neutralized  by  a  series  of  bases  is  always  the  same.  The  strong 
monobasic  acids  (hydrochloric  acid,  hydrobromic  acid,  hydroiodic 
acid,  nitric  acid,  chloric  acid,  bromic  acid,  perchloric  acid,  iodic 
acid),  all  give  off  very  nearly  the  same  amount  of  heat  when  neu- 
tralized by  an  equimolecular  quantity  of  caustic  soda ;  this  amount 
is  very  nearly  138  K.  Among  dibasic  acids,  on  the  contrary,  a 
different  behavior  is  observed.  Although  some,  like  the  monobasic 
acids,  liberate  139  K  for  each  equivalent ;  t  others,  on  the  other 
hand,  liberate  more.  For  instance,  if  increasing  quantities  of  sul- 
phuric acid  are  added  to  a  weight  in  grams  of  sodium  hydroxide 
equivalent  to  one  combining  weight  of  the  hydroxide,  an  evolution 
of  heat  takes  place  until  sufficient  acid  has  been  added  to  form  the 
secondary  sulphate,  Na2  S04 .  This  amounts  to  157  K  for  one  equiv- 

*  The  comma  introduced  between  the  two  portions  of  a  chemical  formula 
indicates  that  the  portions  of  that  formula  separated  are  to  interact  chemi- 
cally to  form  the  complete  substance. 

t  By  equivalent  is  meant  one-half  the  formula  weight,  or  that  proportion 
by  weight  of  the  acid  which  would  contain  one  part  by  weight  of  hydrogen. 


STRONG  AND   WEAK   BASES.  377 

alent  weight  in  grams,  and  to  314  K  for  one  formula  weight  in 
grams,  of  sulphuric  acid.  If  sulphuric  acid  is  further  added,  an 
absorption  of  heat  is  observed  until  a  limit  of  —  33  K  is  reached. 
This  absorption  takes  place  during  formation  of  the  primary  sul- 
phate (see  page  153),  so  that  we  here  observe  the  phenomenon  of  an 
endothermic  reaction  inaugurated  spontaneously.  The  hydroxides 
of  the  alkali  metals  and  of  the  alkaline  earths  form  a  group  of  bases 
which  act  similarly  toward  the  monobasic  acids ;  when  neutralized 
with  hydrochloric  acid  they  give  off,  for  each  equivalent  weight  in 
grams,  about  139  K.  Other  hydroxides,  such  as  aluminium  or  ferric 
hydroxide,  have  much  smaller  heats  of  neutralization.  The  fact 
that  so  much  heat  is  evolved  when  acids  are  brought  in  contact 
with  bases  explains  why  the  reactions  of  neutralization  take  place 
so  readily,  and  why  the  salts  formed  are  so  often  among  the  most 
stable  chemical  bodies.  Of  course,  some  very  weak  acids  are  able 
only  partially  to  neutralize  bases,  and  very  weak  bases  to  partially 
neutralize  acids.  The  heats  of  neutralization  in  these  cases  will  be 
small,  and  the  salts  easily  decomposed  by  the  addition  of  water  or 
of  acids,  or  even  by  slight  warming.  Such  examples  have  been  en- 
countered in  the  study  of  the  chlorides  of  bismuth  and  antimony 
and  in  the  study  of  hydrocyanic  acid.* 

When  we  use  the  expressions  "  strong  "  and  "  weak  "  bases,  it, 
however,  becomes  as  necessary  to  define  the  meaning  of  those  terms 
as  it  was  when  we  used  similar  designations  in  regard  to  acids  (see 
page  141).  But  little  work,  as  compared  with  that  done  in  the 
study  of  acids,  has  been  accomplished  in  regard  to  the  affinity  of 
bases,  yet  we  can  place  the  ratio  between  the  strengths  of  two  bases 
as  being  given  by  the  relative  rapidity  with  which  the  two  are  able 
to  decompose  a  salt  of  a  third,  not  very  pronouncedly  metallic,  sub- 
stance, f  Experiments  in  this  direction  have  shown  that  the  hy- 
droxides of  the  alkali  metals  all  act  about  equally  in  this  respect, 

*  See  pages,  253,  259,  296. 

t  Such  a  substance  is  methylacetate,  CH3  COOCH3 .  In  this  substance 
the  monovalent  radicle  methyl  (page  277)  takes  the  place  of  a  metal.  The 
reaction  between  methylacetate  and  potassium  hydroxide  could  be  expressed 
by  the  following  equation :  — 

CH3  COOCH3  +  KOH  =  CH3  COOK  +  CH3  OH, 
potassium  acetate  and  methyl  hydroxide  (methyl  alcohol)  being  formed. 


378  MASS    ACTION. 

the  alkaline  earths  are  but  little  behind  the  alkalies,  while  ammonia 
develops  a  very  slow  action. 

Our  previous  chemical  study  has  shown  that  chemical  changes 
depend  upon  the  quantity  of  heat  which  is  produced  or  absorbed 
during  the  various  reactions,  and  that  they  also  depend  greatly  upon 
the  temperature  and  other  external  conditions  under  which  these 
reactions  take  place.  Now,  it  is  also  a  matter  of  common  experi- 
ence that  the  mass  of  the  active  chemical  reagents  has  exactly  the 
same  influence  as  the  temperature,  in  such  a  way  that  an  increase 
of  the  mass  *  may  bring  about  the  same  effect  as  a  diminution  of 
temperature,  and  a  diminution  of  the  mass  has  the  same  effect  as 
an  increase  of  temperature,  and  vice  versa.  This  relationship  is 
made  clear  in  reactions  in  which  dissociation  takes  place ;  so,  for  in- 
stance, if  molecules  of  N"204  are  dissociated  by  heat  to  form  NO2, 
the  reunion  of  molecules  of  N02  to  form  N2  04  will  take  place  the 
more  frequently  in  the  unit  of  time,  the  more  often  contact  between 
the  molecules  take  place;  but  this  contact  will  take  place  the  of tener, 
the  greater  the  density  of  the  gas,  i.e.,  the  smaller  the  space  which 
a  given  quantity  has  at  its  disposal ;  the  density  of  a  gas  increases 
with  the  pressure  upon  it,  and  by  this  means  the  active  mass  in 
the  unit  volume  becomes  greater.  From  the  foregoing  it  follows 
that  the  amount  of  N02  present  is  smaller  the  greater  the  density, 
but  it  is  also  true  that  the  amount  of  N02  will  be  smaller  the  lower 
the  temperature.  As  a  result  of  this  and  many  similar  investiga- 
tions, the  law  has  become  well  established  that  the  amount  of  chem- 
ical action  which  a  substance  can  exert  in  any  case  is  proportional 
to  the  active  mass  of  that  substance  which  is  present  in  the  unit  vol- 
ume. Some  of  the  best  illustrations  of  this  law  are  found  in  the 
domain  of  organic  chemistry,  and,  therefore,  are  out  of  place  in 
a  book  of  this  kind  ;  but,  nevertheless,  as  an  understanding  of  some  of 
these  cases  is  necessary  for  the  thorough  comprehension  of  the  de- 
ductions which  are  to  follow,  the  simple  results  of  one  such  inves- 
tigation will  be  detailed.  It  has  been  shown  that  alcohols  can 
react  with  acids  in  much  the  same  way  as  inorganic  hydroxides  do. 
Methyl  alcohol  (methyl  hydroxide)!  when  brought  in  contact  with 
acetic  acid  will  produce  methyl  acetate  and  water,  in  a  manner  simi- 

*  I.e.,  of  the  quantity  of  reagent  contained  in  the  unit  volume. 
j  The  alcohols  are  hydroxides  of  organic  radicles  and,  in  some  respects, 
correspond  to  inorganic  bases.     Thus,  methyl  alcohol  is  methyl  hydroxide; 


DOUBLE    DECOMPOSITION.  379 

lar  to  the  production  of  potassium  acetate  and  water  from  acetic 
acid  and  potassium  hydroxide  ;  the  difference  between  the  two  reac- 
tions being  that  methyl  alcohol  and  acetic  acid  react  slowly  and  (if 
no  provision  is  made  to  remove  the  water  produced)  incompletely, 
while  potassium  hydroxide  and  acetic  acid  neutralize  each  other  at 
once.  Now,  in  such  a  reaction  as  that  which  takes  place  between 
methyl  alcohol  and  acetic  acid,  the  first  change  can  be  represented 
as  follows :  — 

MeOH  +  CH3  COOH  =  CH3  COO  Me  +  H2  0 

Methyl  hydroxide  -\-  Acetic  acid         =  Methyl  acetate    -f  Water ; 

however,  after  a  certain  amount  of  reaction  has  gone  on,  the  mass 
of  the  water  which  is  formed,  becomes  so  great  that  it  begins  to  de- 
compose methylacetate  into  methyl  alcohol  and  acetic  acid :  — 

CH3  COO  Me  +  H2  0  =  CH3  COOH  +  Me  OH, 

so  that  a  point  is  finally  reached  at  which  exactly  as  many  mole- 
cules of  methyl  acetate  are  decomposed,  as  are  formed  in  a  given 
interval  of  time ;  the  entire  system  is  then  in  a  state  of  equilibrium. 
Obviously,  an  increase  in  the  mass  of  water  will  bring  with  it  an 
increased  decomposition  of  acetate,  and  vice  versa. 

Similar  conditions  of  equilibrium  are  encountered  when  two 
salts  in  solution  are  mixed. 

Let  us  suppose  two  substances,  A  B  and  C  D  to  be  in  solution, 
and  let  us  suppose  that  A  B  acts  on  CD  to  produce  two  new  sub- 
stances, A  D  and  B  C.  At  the  first  instant  of  the  reaction  the 
substances  in  solution  can  be  expressed  by  the  following :  — 

AB    CD 
AD    BC. 

These  changes  will  go  on  until  a  point  is  reached  in  which  A  D  and 
B  C  will  be  present  in  such  mass  that  A  D  and  B  C,  reacting  on 
each  other,  will  reproduce  exactly  as  many  molecules  of  A  B  and 
C  D  as  A  B  and  C  D  will  produce  of  A  D  and  B  C.  The  solution  is 
then  in  a  state  of  equilibrium,  which  can  be  disturbed,  however,  by 

ethyl  alcohol  is  ethyl  hydroxide,  etc.  If  we  designate  methyl  by  the  symbol 
Me,  ethyl  by  the  symbol  Et,  then  the  relationship  between  the  structure  of 
alcohols  and  inorganic  bases  is  made  plain  by  the  following  formulae:  — 

Me  OH,  Et  OH,  KOH, 

Methyl  alcohol;     Ethyl  alcohol;     Potassium  hydroxide. 


380  DISSOCIATION    BY   SOLUTION. 

increasing  the  mass  of  one  or  the  other  of  the  constituents.  Let  us 
suppose,  however,  that  A  D,  which  is  produced,  is  either  insoluble 
or  volatile,  it  is  then  removed  from  the  solution  as  fast  as  it  is 
formed,  and  let  us  further  suppose  that  B  C  can  have  no  effect  on 
the  insoluble  or  volatile  substance  produced.  A  familiar  example 
of  such  a  change  would  be  the  action  of  silver  nitrate  on  hydro- 
chloric acid,  by  which  means  silver  chloride,  which  is  insoluble,  and 
nitric  acid  are  produced  :  — 

AgN03  +  H  Cl  =  AgCl  +  HNO». 

In  such  a  case,  then,  as  the  silver  chloride  is  removed  from  the  solu- 
tion as  fast  as  it  is  produced,  no  equilibrium  can  result  until  the 
entire  mass  of  silver  nitrate  has  been  converted  into  silver  chloride. 
The  same  would  be  true  if  the  substance  produced  were  a  gas,  and 
as  a  consequence,  would  as  certainly  be  removed  as  if  it  were  an  in- 
soluble solid.  These  reactions,  known  as  double  decompositions,  are 
among  the  most  common  which  are  to  be  considered  in  the  chemis- 
try of  the  metals.  We  have  already  encountered  a  number  of  them, 
an  instance  being  the  action  of  hydrogen  sulphide  on  soluble  salts 
of  the  heavy  metals  (see  page  100)  and  the  empiric  rule  has  come 
about  as  the  result  of  experience  that  a  complete  double  decomposi- 
tion takes  place  when  an  insoluble  or  volatile  substance  is 
produced.* 

In  discussing  the  determination  of  molecular  weights  by  means 
of  the  lowering  of  the  freezing  point  of  solutions  (see  page  350),  we 
learned  that  if  quantities  of  substances,  so  selected  that  they  are 
proportional  to  their  molecular  weights,  are  dissolved  in  equal 
amounts  of  the  same  solvent,  then  the  lowering  of  the  freezing 
point  of  that  solvent  will  be  the  same  in  each  case ;  when  one 
molecular  weight  of  a  body  is  dissolved  in  one  hundred  molecular 
weights  of  water,  this  depression  amounts  to  0°.62.  This  rule  has, 
however,  not  been  found  to  hold  good  with  many  salt  solutions,  for, 
with  these,  the  depression  is  much  greater  than  would  be  expected ; 
indeed,  with  sodium  chloride  or  potassium  chloride,  it  is  very  nearly 
twice  the  calculated  amount.  A  solution  of  sodium  chloride,  there- 
fore, behaves  as  if  the  quantity  of  salt  represented  by  the  formula 

*  The  laws  -of  mass  action  have  of  late  been  greatly  developed  by  a  num- 
ber of  prominent  chemists ;  their  discussion  is,  however,  out  of  place  in  an  ele- 
mentary text-book.  For  more  detailed  information  the  pupil  can  refer  to 
Ostwald's  Outlines  of  General  Chemistry,  Walker's  translation. 


DISSOCIATION    BY    SOLUTION.  381 

Na  Cl  were  not  one,  but  were  in  reality  two  molecules :  it  seems 
probable,  therefore,  that  the  mere  act  of  solution  dissociates  the 
sodium  chloride  (NaCl)  into  its  ions  (Na-j-  Cl),  and  the  same  is 
true,  although  in  a  lesser  degree,  of  other  salts.  It  has  been 
pointed  out  that  exactly  those  substances  which  (like  sodium  chlo- 
ride) are  electrolytes  (see  page  359),  are  the  ones  which  exhibit  this 
anomalous  behavior  when  in  solution.  Substances  which  are  not 
electrolytes  follow  the  usual  rule.  At  first  glance  it  seems  rather 
startling  that  it  is  those  bodies  which,  owing  to  the  great  energy 
displayed  in  their  formation,  we  have  regarded  as  being  among  the 
most  stable  compounds,  should  be  the  very  ones  which  are  so 
readily  decomposed  on  solution;  but  the  theory  becomes  more 
probable  if  we  remember  that  salts  and  acids  (the  substances  disso- 
ciated by  solution)  enter  into  a  large  number  of  chemical  reactions 
without  difficulty,  and  that,  therefore,  their  constituent  parts  are 
not  so  firmly  united  as  we  would  imagine.  It  must  not  be  thought, 
however,  that  we  can  prove  the  presence  of  free  chlorine  or  free 
sodium  in  a  solution  of  sodium  chloride,  unless  we  have  collected 
the  elements  at  two  electrodes  by  means  of  an  electric  current.  The 
dissociated  ions  themselves,  when  in  solution,  are  possibly  charged 
with  negative  and  positive  electricity,  and  in  this  they  differ  from 
the  free  elements.  According  to  the  theory  which  has  been  out- 
lined, the  neutralization  of  a  base  by  an  acid  in  solution  might  be 
represented  as  follows  (MOH  representing  the  hydroxide  of  any 
metal,  and  HA  any  acid)  :  *  — 

M  +  OH       +H+A  =         M  +  A   +  H2  0 

Dissociated  base  +  Dissociated  acid  =  Dissociated  salt  +  water. 
A  glance  at  the  above  equation  will  demonstrate  that  the  change 
in  condition  in  the  entire  system  consists  simply  in  the  formation 
of  water  from  its  ions,  H  and  OH,  and  this  formation  is  obviously 
independent  of  the  nature  of  the  acid  (AH)  and  of  the  base  (MOH). 
This  view  explains  fully  the  fact  that  the  heat  of  neutralization  is 
independent  of  the  nature  of  the  base  and  of  the  nature  of  the  acid 
(see  page  376). 

Now,  it  has  been  demonstrated  that  all  electrolytes  which  man- 

*  See  Ostwald;  Zeitschrift  fur  Physikal.  Chemie;  7,  423;  and  also  Meyer; 
Grundziige  der  Theoret.  Cliemie;  p.  360  and  following 


382  DISSOCIATION    BY    SOLUTION. 

ifest  the  abnormal  lowering  of  the  freezing  point  which  has  been 
mentioned  (which  are  therefore  partially  dissociated  on  solution) 
conduct  electricity  the  more  readily  the  greater  the  dissociation. 
Dissociation  of  an  electrolyte  seems,  therefore,  to  be  essential  for 
its  conducting  electricity,  and  it  is  probably  for  this  reason  that 
pure  water  and  pure  liquid  hydrochloric  acid,  for  example;  do  not 
conduct  electricity,  but  that  a  solution  of  hydrochloric  acid  in  water 
does.  At  high  temperatures,  however,  many  pure  substances  are 
electrolytes  (pure  sodium  chloride,  pure  potassium  chloride,  for 
example)  ;  but  then,  at  high  temperatures,  the  tendency  toward 
decomposition  is  already  a  great  one. 

With  the  hypothesis  of  dissociation  in  view,  it  is  not  necessary 
to  assume  that,  as  has  been  stated  above,  two  salts,  on  mixing  their 
solutions,  in  reality  yield  four  salts  by  double  decomposition.  Let 
us  take  the  following  case  for  an  example.  Suppose  we  have  a 
solution  of  sodium  chloride ;  this  will  contain  free  ions  of  sodium 
and  of  chlorine,  and  now  let  us  add  to  this  a  solution  of  potassium 
nitrate  (ions,  K  and  N03)  ;  obviously,  if  dissociation  is  nearly  com- 
plete, such  a  solution  will  be  identical  in  every  respect  with  the 
one  produced  by  mixing  sodium  nitrate  with  potassium  chloride,  a 
result  which  is  identical  with  that  which  would  be  arrived  at  were 
double  decomposition  (see  page  379)  to  take  place.  A  more  com- 
plete discussion  of  these  topics  belongs  in  a  text-book  more  espe- 
cially devoted  to  physical  chemistry.  It  will  probably  be  shown,  in 
time,  that  all  reactions  which  take  place  at  ordinary  temperatures 
and  in  solution,  are,  of  necessity,  preceded  by  dissociation. 


THE   ALKALI   METALS.  383 


CHAPTER   LII. 

THE   ALKALI   METALS. 

Lithium, ;  symbol,  Li ;   atomic  weight,  7.02. 
Sodium  ;  symbol,  Na ;  atomic  weight,  23.05. 
Potassium. ;  symbol,  K  ;  atomic  weight,  39.11. 
Rubidium  ;  symbol,  Eb ;  atomic  weight,  85.5. 
Caesium, ;  symbol,  Cs ;  atomic  weight,  132.9. 

THE  alkali  metals  are  the  chemical  opposites  of  the  halogens ; 
the  latter  are  the  most  negative  (not-metallic)  elements  with  wtdch 
we  are  acquainted,  while  the  former  are  the  most  posit^er^  Tme 
metallic  properties  of  the  alkalies  increase  with  increasing  atomic 
weight,  just  as  the  not-metallic  properties  in  the  halogen  group 
diminish  from  fluorine  to  iodine ;  these  changes  can,  perhaps,  best 
be  studied  by  a  comparison  of  the  readiness  with  which  the  indi- 
vidual members  of  both  families  decompose  water.  This  decompo- 
sition is  of  an  exactly  opposite  character,  accordingly  as  the  element 
in  question  is  a  member  of  the  halogen  or  of  the  alkali  group.  As 
a  reference  to  page  75  will  recall,  the  halogens  decompose  water, 
liberating  oxygen,  and  they  do  this  the  more  readily  the  more  neg- 
ative they  are;  on  the  other  hand,  the  alkalies  decompose  water, 
liberating  hydrogen,  and  this  reaction  takes  place  the  more  readily 
the  more  positive  the  metal  in  question  is.  Fluorine,  when  brought 
in  contact  with  water,  instantly  forms  hydrofluoric  acid,  and  sets 
free  oxygen,  even  in  the  absence  of  light;  chlorine  does  so  only 
when  the  solution  of  that  gas  is  placed  in  the  sunlight ;  bromine 
enters  into  this  reaction  more  slowly  than  chlorine,  while  iodine 
has  no  effect.  In  the  case  of  the  alkali  metals,  lithium  decomposes 
water,  forming  lithium  hydroxide  and  hydrogen,  but  the  metal 
does  this  quietly,  without  itself  melting  or  without  generating 
sufficient  heat  to  cause  the  hydrogen  to  take  fire ;  sodium  attacks 
water  energetically ;  the  metal  is  heated  to  its  melting  point,  but 
the  hydrogen  which  is  being  evolved  does  not  burst  into  flame  ; 
potassium  melts,  while  sufficient  heat  is  developed  to  ignite  the 


384 


THE   ALKALI   METALS. 


hydrogen ;  while  both  rubidium  and  caesium  enter  into  the  decom- 
position with  explosive  violence.  The  above  relationships  are  more 
clearly  shown  by  the  following  table  :  — 


THE  HALOGENS. 


DECOMPOSITION  OF  WATER. 

NOT-METALLIC  CHARACTER. 

ELEMENTS. 

ATOMIC  WEIGHTS. 

F, 

19. 

- 

Cl, 

35.45 

• 

Br, 

79.95 

Liberating  oxygen. 

I, 

126.85 

THE  ALKALI  METALS. 


DECOMPOSITION  OF  WATEB. 

METALLIC  CHARACTER. 

ELEMENTS. 

ATOMIC  WEIGHTS. 

M 

H 

Li, 

7.02 

^r 

y 

Na, 

23.05 

vL 

vL 

K, 

39.11 

T 

T 

Kb, 

85.5 

Liberating  hydrogen. 

Cs, 

132.9 

The  decomposition  of  water  takes  place  with  increasing  readi- 
ness in  the  direction  of  the  arrows ;  the  not-metallic  character  in 
the  halogen  family,  the  metallic  character  in  the  alkali  family, 
increase  in  the  direction  of  the  arrows.  Of  course,  the  halogens 
and  the  alkali  metals  unite  with  the  greatest  energy;  the  latter 
burn  in  an  atmosphere  of  the  former  with  a  brilliant  light,  while 
most  stable  halides  are  produced. 

The  alkali  metals  are  soft,  malleable,  and  ductile,  and  possess  a 
brilliant  metallic  lustre.  When  they  are  exposed  to  the  air,  they 
almost  instantly  become  coated  with  a  layer  of  oxide  ;  this  oxide 
absorbs  moisture  and  carbon  dioxide,  and  is  soon  converted  into  a 
mixture  of  the  carbonate  and  hydroxide ;  as  a  consequence  of  these 
changes,  pieces  of  the  alkali  metals  which  are  in  contact  with  the 
atmosphere  become  the  centres  of  small  pools  of  deliquescent  hy- 
droxide. Because  of  this  capacity  for  oxidation,  the  metals  are 
kept  under  some  liquid  hydrocarbon,  such  as  petroleum  naphtha. 

Contrary  to  the  rule  observed  with  not-metallic  elements  (which 
form  the  opposite  extremity  of  the  periods),  the  metals  comprising 
a  family  which  finds  its  place  at  the  beginning  of  these  periods, 


ALKALI    METALS  ;    OXIDES.  385 

show  a  diminution  of  their  melting  points,  with  an  increase  in  their 
atomic  weights ;  this  change  is  evident  from  the  following  table  :  — 

Li,  specific  gravity,  .589  atomic  volume,  11.9  melting  point,  180° 

Na,      "             "  .972  "           "        23.6         "          "  95°.6 

K,         "              "  .865  "           "        45.2          "           "  62°.5 

Kb,       "              "  1.52  *'           "        56.2          "           "  38°.5 

Cs,        "              "  1.88  "           "        70.6          "          "  26°   . 

The  alkaline  oxides  are  soluble  in  water,  and  when  so  dissolved 
they  produce  the  corresponding  hydroxides ;  from  this  it  follows 
that  the  former  cannot  be  made  excepting  by  processes  during 
which  water  is  rigidly  excluded.  Each  of  the  elements  can  be 
burned  in  oxygen  to  form  the  corresponding  oxide,  which  is  con- 
structed according  to  the  general  formula  M2  0  ;  the  reaction  by 
which  the  oxides  change  to  the  hydroxides  can  be  represented  by 
the  following  equation :  — 

M2  0  +  Ho  0  =  2  MOH. 

During  the  change  from  oxide  to  hydroxide,  and  during  the  subse- 
quent solution,  a  large  amount  of  heat  is  given  off ;  for  instance,  in 
the  case  of  Na20,  this  amounts  to  550  K.  The  hydroxides  are, 
without  exception,  soluble  ;  this  solubility  increases  with  the  metal- 
lic nature,  and,  hence,  with  the  atomic  weights  and  volumes  of  the 
alkali  metals,  the  more  positive  the  metal,  therefore,  the  more  solu- 
ble the  hydroxide.  The  rule  is  so  far  without  exception  that,  in 
no  other  family  of  metals  do  we  encounter  hydroxides  which  are 
readily  soluble  in  water.*  If  we  examine  the  next  family  to  the 
right  of  the  alkalies  in  the  periodic  system,  we  shall  discover  that 
the  hydroxides  of  the  two  members  with  the  smallest  atomic 
weights  (beryllium  and  magnesium)  are  insoluble,!  while  those  of 
calcium,  strontium,  and  barium,  although  in  no  case  so  readily  sol- 
uble as  the  hydroxides  of  the  alkali  metals,  have  their  solubility 
increased  with  increasing  atomic  weight,  so  that  the  same  rule,  hold- 
ing good  with  the  alkali  metals,  appertains  to  this  family  also. 
With  the  families  mentioned  (the  alkalies  and  the  alkaline  earths), 
the  list  of  soluble  hydroxides  of  the  metals  is  practically  exhausted, 
for  by  far  the  greater  number  of  hydroxides  of  the  purely  metallic 
elements  are  insoluble  in  pure  water. 

*  The  hydroxide  of  lithium,  in  the  alkali  family,  is  not  soluble  with  great 
readiness.     Lithium  is  also  the  least  positive  of  the  alkali  metals. 
t  That  of  magnesium  is  very  nearly  insoluble. 


386  ALKALI  METALS;    HYDROXIDES. 

/ 

The  hydroxides  of  the  alkali  metals  (the  caustic  alkalies)  can  be 
prepared  by  the  action  of  the  lespective  metals  on  water. 

M  -f  HOH  =  MOH  +  H, 

or  by  covering  slaked  lime  with  a  solution  of  an  alkaline  carbonate, 
allowing  the  mixture  to  stand,  and  then  filtering,  when  the  following 
reaction  has  taken  place  •  — 

M2  C03  +  Ca  (OH  )2  =  Ca  C03   +  2  MOH, 

Soluble.         Soluble.  Insoluble.       Soluble. 

The  filtered  liquid  is  evaporated,  at  first  in  porcelain,  and  finally  in 
iron  or  silver  dishes.*  All  of  the  alkaline  hydroxides  can  be  fused 
without  decomposing  into  the  corresponding  oxide  and  water,  and 
the  solution  of  any  one  of  them  absorbs  carbon  dioxide  when  ex- 
posed to  the  air,  so  that,  if  it  is  to  be  kept  for  any  length  of  time, 
it  must  be  placed  in  closed  flasks.  The  solutions  are  strongly  alka- 
line in  taste  and  in  reaction  toward  litmus,  and  neutralize  acids  with 
the  greatest  readiness. 

The  sulphides  of  the  alkali  metals  bear  a  great  resemblance  to 
the  oxides  ;  they  are  all  quite  soluble  in  water,  and  are  formed,  as 
a  general  rule,  by  reduction  of  the  corresponding  sulphates  by  heat- 
ing with  charcoal :  — 

M2  S04  +  4  C  =  M2  S  +  4  CO, 

or  by  addition  of  hydrogen  sulphide  to  a  solution  of  the  correspond- 
ing hydroxide,  by  which  means  the  sulphhydrate  is  produced :  — 

MOH  -f  H2  S  =  MSH  -f  H2  0. 

An  addition  of  an  equal  amount  of  hydroxide  to  the  sulphhydrate 
then  produces  the  sulphide  :  — 

MSH  +  MOH  =  M2  S  +  H2  0. 

(In  this  case  H2  S  acts  like  a  dibasic  acid ;  see  page  139.) 
The  solutions  of  alkaline  sulphides  are  able  to  dissolve  sulphur  to 
form  so-called   polysulphides    (see   page    155,  foot-note).      In  the 
cases  of  sodium  and  potassium,  polysulphides  having  the  following 
formulae  have  been  isolated  :  — 

Na2S2,        Na2S4,  K2S2,        K2S4, 

Na2S3,         Na2S5,  K2S3,         K2S5. 

*  Platinum  vessels  must  not  be  heated  with  concentrated  solutions  of 
alkaline  hydroxides,  for  they  are  readily  attacked  by  caustic  alkalies. 


ALKALI  METALS;  SULPHIDES.  387 

All  of  these,  on  addition  of  acids,  form  the  corresponding  salt, 
hydrogen  sulphide,  and  sulphur,  for  instance  :  — 

K2  S5  +  2  H  Cl  =  2  K  Cl  +  H2  S  +  4  S. 

When  exposed  to  the  air  they  are  oxidized,  forming  salts  of 
thiosulphuric  acid,  while  sulphur  is  liberated :  — 

K2S2  +  30=K2S203, 

K2  S3  +  3  0  =  K2  S2  03  +  S, 

K2  S4  +  3  0  =  K2  S2  03  +  2  S, 

K2  S5  +  3  0  =  K2  S2  03  +  3  S, 

The  formulae  of  these  sulphides  can,  perhaps,  be  best  explained  on 
the  supposition  that  they  are  salts  of  thio  or  sulpho  acids ;  so  that 
the  action  of  sulphur  on  the  monosulphides  would  be  analogous  to 
that  of  oxygen  on  the  same  substances.  For  example,  the  sulphide 
of  potassium  would  be  oxidized  to  the  sulphate :  — 

K2S  +  40  =  K2S04," 
while  the  sulphide  would  be  sulphurized  to  the  pentasulphide :  — 

K2  S  +  4  S  =  K2  S5 . 

According  to  this  theory,  the  pentasulphide  would  be  the  potassium 
salt  of  dithio-disulpho  sulphuric  acid,  or :  — 

SK 

S 
S 
SK. 

This  theory  is,  however,  sustained  only  by  the  fact  that  the 
sulphides  of  the  alkalies  can  take  up  no  more  than  four  atoms  of 
sulphur.  The  acids  corresponding  to  these  sulphides  have  never 
been  isolated.* 

The  sulphhydrates  of  the  alkali  metals,  corresponding  to  the 
hydroxides  in  formula,  are  soluble  in  water,  and  are  produced  by 
the  action  of  hydrogen  sulphide  on  the  hydroxides :  — 

KOH  +  H2  S  =  KSH  +  H2  0. 

Both  the  sulphides  and  the  sulphhydrates  are  bases  ;  with  acids  they 
form  salts  and  hydrogen  sulphide  (the  oxides  and  hydroxides  form 
salts  and  water  ;  see  page  78)  :  - 

*  This  theory  as  to  the  constitution  of  the  polysulphides  was  first  proposed 
by  Drechsel  (Journal  fur  Prakt.  Chemie;  [2]  4,  20).  It  has  also  been  adopted 
by  Remsen  in  his  Inorganic  Chemistry,  p.  207. 


388  ALKALI  METALS;   HALIDES. 

K20  +2HC1  =  2KC1  +  HS0 
K2S  +  2HC1  =  2KC1  +  H2S 
KOH  +  H  Cl  =  K  Cl  +  H2  0 
KSH  -f  H  Cl  =  K  Cl  +  H2  S. 

The  halogen  derivatives  of  the  alkalies  are  extremely  stable,  and 
are,  throughout,  soluble  in  water.  Of  these  compounds,  the  most 
common  in  occurrence  are  the  chlorides,  of  which  sodium  chloride 
is  by  far  the  most  often  met  with.  Sodium  chloride,  or  common 
salt,  occurs  in  extensive  beds  in  rocks  of  various  ages,  associated 
with  gypsum,  calcite,  clay,  and  sandstone.  It  frequently  occurs  in 
solution  in  salt  springs,  and  is  always  found  in  the  sea,  of  which 
it  forms  2.5  per  cent.  The  salt  of  commerce  is  often  obtained  by 
evaporating  sea  water  in  lagoons  by  means  of  the  heat  of  the  sun, 
as  is  done  in  France.  Lime,  gypsum,  and  ferric  hydrate  separate 
at  first;  afterward  the  salt  begins  to  crystallize  and  can  be  raked 
out;  at  last  there  is  left  a  mother  liquor  which  contains  sodium 
chloride,  magnesium  chloride,  potassium  chloride,  and  magnesium 
sulphate.  In  many  cases,  where  the  brine  obtained  from  salt 
springs  is  evaporated,  the  mother  liquors  contain  the  bromides  and 
iodides  of  the  alkalies.  The  simple  halides  of  the  alkali  metals 
crystallize  in  the  regular  system,  most  frequently  in  cubes.  As  we 
have  seen  (page  337),  the  halides  of  the  alkali  metals  have  a 
tendency  to  crystallize  with  other  halides  in  the  form  of  double 
salts,  in  which  the  alkaline  halide  presumably  plays  the  part  of  the 
base.  A  number  of  these  double  halides  occur  as  natural  minerals. 
Among  the  most  important  of  them  is  cryolite,  Al  F3 ,  3  Na  F  (see 
page  333).  None  of  the  bromides  or  iodides  of  the  alkalies  occur 
as  crystalline  mineral  individuals. 

As  would  be  expected,  the  heat  of  formation  of  the  alkaline 
halides  increases  with  the  increasing  metallic  nature  of  the  alkali 
metal  forming  such  a  halide.  This  relationship  is  readily  seen  from 
the  following  table  :  — 

HEAT  OF  FORMATION  OF  THE  CHLORIDES. 

Li  Cl,  938  K, 
Na  Cl,  976  K, 
K  Cl,  1043  K. 

A  distinctive  feature  of  the  chemistry  of  the  alkali  metals  lies 
in  the  fact  that  the  salts  of  these  metals  are  almost  without  excep- 


SODIUM   CARBONATE.  389 

tion  soluble  in  water ;  they  can  be  produced  by  neutralizing  the 
hydroxide  solutions  with  the  various  acids. 

The  carbonates  of  the  alkalies  are  soluble  in  water,  differing  in 
this  way  from  those  of  the  alkaline  earths ;  however,  being  the  car- 
bonates of  the  most  positive  metals,  they  are  not  decomposed  by 
heat,  as  are  the  same  salts  of  all  other  elements ;  the  carbonate  of 
the  least  metallic  element  of  the  alkali  family  ( lithium)  is  also  the 
least  soluble  in  water ;  the  solubility  of  the  carbonates,  as  we  pass 
from  member  to  member  in  this  group,  increases  with  the  increase 
in  the  metallic  character,  and  hence  of  the  atomic  weight  of  the 
alkali  metal  producing  the  salt.  The  most  important  carbonates 
are  those  of  sodium  and  of  potassium,  and,  as  sodium  carbonate 
(common  soda)  is  of  great  commercial  importance,  it  is  advisable  to 
enter  into  a  brief  description  of  the  process  of  its  manufacture. 

Sodium  carbonate  occurs  in  some  mineral  waters  (Karlsbad)  and 
as  a  remainder  after  evaporating  the  water  of  alkaline  lakes,  it 
is  also  a  constituent  of  the  ashes  of  sea-plants,*  and  it  was  from 
these  that  the  soda  of  commerce  was  made  until  the  end  of  the  last 
century.  During  the  period  of  the  French  Revolution,  a  large 
reward  was  offered  for  the  discovery  of  a  process  by  means  of  which 
sodium  carbonate  could  be  prepared  from  the  chloride,  as  the  latter 
substance  was  a  product  which  was  both  cheap  and  easily  purified. 
Nicholas  LeBlanc,f  owing  to  this  inducement,  discovered  a  process 
which  has  been  used  with  but  little  modification  up  to  the  present 
day.  The  chemical  changes  upon  which  this  method  depends  are 
as  follows  :  — 

Sodium  chloride  is  treated  with  sulphuric  acid)  when  hydrochlo- 
ric acid  and  primary  sodium  sulphate  are  formed :  — 

Na  Cl  +  H2  S04  =  Na  HS04  +  H  01. 

The  primary  sulphate  is  then  heated  with  sodium  chloride,  produ- 
cing hydrochloric  acid  and  the  secondary  sulphate  :  — 

Na  Cl  +  Na  HS04  =  Na2  S04  +  H  Cl. 

The  hydrochloric  acid  which  passes  off  is  absorbed  by  water  and  is 
used  as  ordinary  commercial  hydrochloric  acid. 

The  sodium  sulphate  is  next  converted  into  crude  soda  by  heating 

*  The  ashes  of  land  plants  consist  mainly  of  potassium  carbonate, 
t  Physician  to  the  Due  <T  Orleans. 


390  SODIUM   CARBONATE. 

with  anthracite  coal  and  chalk  (calcium  carbonate),  the  tempera- 
ture reaching  1000°.  Despite  the  extended  attention  given  to  the 
soda  manufacture  in  the  last  few  years,  the  chemical  processes 
taking  place  have  not  been  definitely  settled,  yet  the  following  are 
most  frequently  accepted  as  nearly  correct  :  Sodium  sulphate  first  is 
reduced  to  sodium  sulphide  by  the  coal  :  — 

Na2  S04  +  2  C  =  Na2  S  +  2  C  02  ; 

and  the  sodium  sulphide,  together  with  chalk,  then  changes  to 
sodium  carbonate  and  calcium  sulphide  :  — 

Na2  S  -f>  Ca  C03  =  Na2  C03  +  Ca  S. 

The  crude  soda  is  extracted  with  water  and  the  liquors  are  evapo- 
rated, when  tolerably  pure  sodium  carbonate  crystallizes.  This,  when 
slowly  crystallized  from  water,  separates  with  ten  molecules  of  that 
substance  and  forms  crystals  of  commercial  soda,  Na2  C03  -f-  10  H2  0  . 
The  latter  effloresce  when  in  contact  with  the  air,  losing  9  molecules 
of  water  and  changing  into  a  powder,  which  has  the  composition 


A  modern  process,  known  as  the  ammonia-soda  piwcess,  has  of  late 
succeeded  to  a  large  extent  in  taking  the  place  of  the  older  method. 
This  late  improvement  depends  on  the  fact  that  primary  sodium 
carbonate  is  not  very  readily  soluble  in  cold  water.  Ammonia 
solution,  saturated  with  an  excess  of  carbon  dioxide,  contains  pri- 
mary ammonium  carbonate  (NH4)  HC03,  and  when  this  is  added  to 
a  solution  of  common  salt  the  following  change  takes  place  :  - 

(  NH4  )  HC03  +  Na  Cl  =  Na  HC03  +  NH4  Cl  ; 

the  primary  sodium  carbonate  (sodium  bicarbonate)  separates  as  a 
crystalline  powder,  and  the  latter,  when  heated,  gives  off  water  and 
carbon  dioxide,  leaving  the  secondary  carbonate  (see  page  291)  :  — 

2NaHC03  =  Na2CO  +  H20  +  C02. 

The  nitrates  of  sodium  and  of  potassium  are  of  importance  in  the 
manufacture  of  gunpowder.  Sodium  nitrate  is  found  in  the  north- 
ern part  of  Chile,*  where  it  occurs  in  extensive  deposits,  accom- 
panied by  sodium  chloride  and  other  salts,  the  presence  of  which 
seems  to  indicate  that  the  formation  of  the  nitrate  is  due  to  the 
decay  of  marine  plants,  the  occurrence  of  these  deposits  in  this 
place  being  one  of  the  proofs  of  the  theory  that  this  portion  of 

*  Province  of  Tarapaca  ;  the  nitrate  is  called  Chile  saltpetre. 


SODIUM   AND   POTASSIUM   NITRATES.  391 

South  America  was  at  one  time  submerged.  Sodium  nitrate  is  puri- 
fied by  washing  with  water  and  recrystallization  ;  the  mother  liquors, 
which  are  left,  contain  considerable  quantities  of  the  iodides  and  are 
used  in  the  manufacture  of  iodine.  Unfortunately,  sodium  nitrate, 
because  it  is  hygroscopic,  cannot  be  used  in  the  preparation  of  gun- 
powder ;  so  that,  preliminary  to  the  production  of  that  explosive,  the 
nitrate  of  sodium  must  be  converted  into  nitrate  of  potassium.  This 
is  accomplished  by  treating  a  saturated  solution  of  sodium  nitrate, 
at  boiling  heat,  with  a  solution  of  potassium  chloride,  when  a  double 
decomposition,  accompanied  by  the  formation  of  the  less  soluble 
chloride  of  sodium,  takes  place :  — 

Na  N03  +  K  Cl  =  Na  Cl  +  KN03 , 

The  solution  of  potassium  nitrate  is  filtered  and  allowed  to  crystal- 
lize. Potassium  nitrate  occurs  as  a  mineral  deposit  in  many  places, 
where  nitrogenous  organic  matter  is  decaying  in  the  presence  of  pot- 
ash (see  page  204) ;  localities  in  which  this  natural  production  of 
potassium  nitrate  assumes  commercial  importance  are  found  in  Spain, 
Egypt,  Peru,  and  especially  India,  from  which  latter  country  potas- 
sium nitrate  is  exported  in  considerable  quantities. 

The  use  of  potassium  nitrate  in  the  manufacture  of  gunpowder 
depends  on  its  oxidizing  powers.  When  potassium  nitrate  is  mixed 
with  charcoal  and  ignited,  the  following  reaction  takes  place. 

2  KN03  +  3  C  =C02  +  CO  +  K2  C03  +  2  N ; 

the  carbon  dioxide,  which  is  left  in  combination  as  potassium  carbon- 
ate, can  be  further  liberated  by  the  previous  addition  of  sulphur :  — 

2  KN03  +  3  C  +  S  ==  3  C02  +  2  N  +  K2  S. 

The  formation  of  such  a  large  amount  of  gaseous  material  from  the 
small  volume  of  solid  causes  the  explosion.  The  reaction  given  is, 
however,  only  approximately  correct,  for  other  changes,  not  defi- 
nitely understood,  also  take  place. 

The  other  salts  of  the  alkalies  have  been  sufficiently  mentioned 
in  the  course  of  the  chapters  in  which  the  various  acids  have  been 
discussed.  The  following  table  will  make  the  formulae  and  solubility 
of  some  of  the  salts  most  apparent :  — 

ALKALI   METALS. 

OXIDES,  M2  O,  converted  to  the  hydroxides  by  addition  of  water. 
HYDROXIDES,  MOH,  soluble  in  water;  least  soluble  is  Li  OH;  solubility 
increases  with  increasing  atomic  weight  of  the  alkali  metal. 


392  SPECTROSCOPE. 

SULPHIDES,   M2  S,  soluble,  probably   converted    into   MSH  +  MOH   by 

addition  of  water. 

SULPHHYDRATES,  MSH,  soluble  in  water;  MOH  +  H2  S  =  MSH  +  H2  O. 
CARBONATES,  M2CO3  and  MHCO3 ,  soluble  in  water;  least  soluble  is 

Li2  COg  ;  solubility  increases  with  increasing  atomic  weight. 
NITRATES,  MNO3 ,  soluble  in  water,  change  to  nitrites  and  oxygen  when 

heated.     (See  page  208.) 

SULPHATES,  M2  SO*  and  MHSO4  .     Soluble  in  water. 
PHOSPHATES,  MH2PO4,  M2HPO4,  M3PO4,  tertiary  phosphates  change 

to  secondary  phosphate  and  hydroxide  of  alkali  metals  on  addition  of 

water.     (See  page  230. ) 
SILICATES,  M2SiO3,  soluble  in  water. 
POTASSIUM  PERCHLORATE,  KC1O4,  potassium  fluosilicate,  K2SiF6,  are 

soluble  with  difficulty,  and  hence  are  precipitated  from  solutions  of 

potassium  salts  by  addition  of  the  corresponding  acids,  sodium  pyro- 

antimonate,  Na2  H2  Sb2  O7 ;  insoluble  in  cold  water. 

The  salts  of  ammonium  correspond  entirely  to  the  salts  of  potas- 
sium, and  hence  are  frequently  discussed  in  connection  with  the 
alkali  metals.  Their  nature  has  been  sufficiently  explained  on  page 
188  and  sub. 

Detection  of  metals  by  means  of  the  spectroscope. 

The  readiest  means  for  the  detection  of  the  various  alkali  metals 
is  by  means  of  the  spectroscope  ;  indeed,  caesium  and  rubidium 
were  not  known  to  exist  until  the  examination,  by  the  spectroscope, 
of  the  residues  left  by  evaporation  of  certain  mineral  waters  re- 
vealed their  presence.  The  principle  upon  which  the  use  of  the 
the  spectroscope  depends  is  as  follows :  — 

Light  which  contains  waves  of  only  one  wave-length  is  mono- 
chromatic (homogeneous).  When  a  ray  of  such  light,  passing 
through  the  air,  comes  in  contact  with  a  transparent  medium  of 
greater  density,  it  is  refracted  toward  a  line  normal  to  the  surface 
of  the  latter,  and  the  smaller  the  wave-length,  the  greater  is  the 
refraction ;  for  each  wave-length  there  is  a  corresponding  index  of 
refraction,  provided  the  media  through  which  each  kind  of  light 
passes  remain  the  same.  If  a  ray  of  light  contains  waves  of  various 
lengths,  then  each  kind  of  wave  will  be  refracted  according  to  its 
refractive  index,  so  that  the  whole  will  be  separated  into  as  many 
monochromatic  rays  as  it  contained  different  wave-lengths.  Such 
a  ray  of  light,  falling  from  a  narrow  slit  upon  a  prism,  the  edge  of 
which  is  parallel  to  the  slit,  produces  a  series  of  parallel  images  of 
the  opening;  and  if  the  light  consists- of  all  of  the  colors  between 


THE    SPECTROSCOPE. 


393 


two  determined  extremes,  the  image  obtained  appears  as  a  continu- 
ous spectrum,  produced  by  a  number  of  different  colored  images  of 
the  slit,  which  merge  the  one  into  the  other.  Light  which  is 
emitted  by  the  sun  or  by  white-hot  bodies,  contains  an  infinite 
number  of  waves  which  differ  in  length ;  this  light  is  white ;  when 
passed  through  a  prism  it  gives  a  spectrum  containing  all  the  colors 
of  the  rainbow,  beginning  with  red  and  passing  through  the  various 
modifications  of  color  (orange,  yellow,  green,  blue)  to  violet,  at  the 
opposite  extremity. 


Fig.  12. 

The  spectroscope  (Fig.  12)  consists  of  a  telescope  (A)  which  throws 
parallel  rays  of  light  admitted  through  a  small  slit  at  (S)  upon  a  prism  (P); 
the  spectrum  formed  is  observed  by  the  telescope  at  B,  which  is  so  focussed  as 
to  give  a  sharp  image  of  the  same;  at  the  same  time  a  mirrored  image  of  a 
millimeter  scale,  photographed  and  placed  in  C,  is  so  reflected  as  to  be  visible 
above  the  spectrum  when  the  observer  glances  through  B.  The  manner  of 
opening  and  closing  the  slit  by  means  of  the  thumb  screw  (e)  is  shown  by  the 
small  figure  (d). 

The  spectrum  of  a  white-hot  solid,  when  so  observed,  is  contin- 
uous, but  this  is  not  the  case  with,  glowing  gases.  These,  when 
examined  by  the  spectroscope,  show  a  number  of  bright,  colored 
lines  upon  a  black  or  nearly  black  background.  The  colors  and 
relative  position  of  the  lines  are  definite  ones  for  each  individual 
glowing  gas,  and  are  always  different  for  gases  of  differing  chem- 
ical composition.  The  reason  for  the  appearance  of  these  lines  is 
•that  the  glowing  gases  emit  light  only  of  certain  determined  wave- 
lengths, the  varieties,  and  hence  the  colors,  of  which  are  generally 


394 


SPECTRUM   ANALYSIS. 


few  in  number ;  as  a  consequence,  light  of  each  wave-length,  being 
refracted  according  to  its  index  of  refraction,  appears  in  a  different 
place  on  the  spectrum  as  a  sharp  line  of  the  color  belonging  to  that 
particular  wave-length. 

The  flame  of  a  Bunsen  burner  is  not  luminous,  but  if  a  platinum 
wire  is  placed  in  it,  the  latter  becomes  heated  and  emits  a  white 
light.  If,  now,  the  wire  is  coiled  as  in  Fig.  13,  and  is  (after  moist- 
ening with  a  little  hydrochloric  acid) 
dipped  into  a  little  sodium  chloride, 
the  adhering  salt,  when  brought  Fig' 13' 

into  the  flame,  will  vaporize  and  will  emit  a  pure  yellow  light.  In 
the  same  way,  the  light  emitted  by  potassium  salts  will  be  violet ;  * 
by  lithium  or  strontium,  red  ;  by  copper,  barium,  or  thallium,  green  ; 
by  zinc,  blue,  etc.f  If  the  yellow  sodium  light,  obtained  as  above, 
is  observed  by  means  of  a  spectroscope,  a  bright  yellow  line  on  a 
dark  background  is  seen.  The  position  of  this  yellow  line  corre- 
sponds to  the  position  occupied  by  yellow  in  the  continuous  spec- 
trum. Lithium  will  show  a  red  line  and  a  less  marked  yellow  one, 

potassium  a  red  line  and  a  blue 
one.  In  fact,  each  individual  metal 
displays  characteristic  lines  in  defi- 
nite parts  of  the  spectrum,  while 
the  lines  of  no  two  metals  corre- 
spond exactly.  In  order,  then,  to 
discover  the  presence  of  any  metal 
or  metals  in  a  mixture,  it  is  only 
necessary  to  place  volatile  com- 
pounds of  those  elements  in  the 
not-luminous  flame  and  then  to  ob- 
The  minutest  traces  of  the  metals 
in  question  can  be  detected  in  this  way. 

If  a  not-luminous  flame  is  placed  in  front  of  a  luminous  back- 
ground which  emits  a  white  light,  and  if  then  some  sodium  salt  is 

*  Best  seen  when  the  potassium  flame  is  observed  through  a  piece  of  blue 
glass.  When  sodium  is  also  present,  the  blue  absorbs  the  yellow  rays,  while 
permitting  the  violet  ones  to  pass  through. 

t  The  contrasting  colors  of  the  various  flames  can  be  best  observed  by 
using  a  lamp  with  four  or  five  burners  (Fig.  14)  and,  after  fixing  platinum 
wires  in  the  stands,  as  shown  in  the  figure,  bringing  the  entire  number  simul- 
taneously into  the  lighted  burners. 


Fig.  14. 

serve  the  spectrum  produced. 


SPECTRUM   ANALYSIS.  395 

volatilized  in  this  not-luminous  flame,  the  spectroscope  will  show  a 
continuous  spectrum,  due  to  the  white  light,  with  this  difference, 
however,  that  in  the  place  where  the  yellow  band  of  light  belonging 
to  the  sodium  spectrum  usually  occurs,  there  is  now  seen  a  black 
band.  This  phenomenon  is  due  to  the  fact  that  when  a  ray  of 
white  light  containing  all  colors  is  passed  through  a  glowing  gas 
which  emits  light  rays  only  of  certain  definite  colors,  this  glowing 
gas  is  able  to  absorb  from  the  white  light  the  rays  of  exactly  the 
same  color  as  those  which  itself  emits.  It  follows  that,  when 
white  light  is  passed  through  the  glowing  vapor  of  a  potassium 
compound,  there  will  appear  (on  the  continuous  spectrum)  a  dark  line 
in  the  red  and  one  in  the  blue ;  in  the  case  of  lithium,  a  dark  line 
in  the  red  and  one  in  the  yellow,  etc.  Such  spectra  are  called  ab- 
sorption spectra.  The  spectra  of  the  sun  and  of  the  fixed  stars  are 
not  perfectly  continuous,  but  are  traversed  by  a  series  of  fine,  dark 
lines  which  have  been  proven  to  correspond  exactly  to  the  absorp- 
tion spectra  of  the  glowing  vapors  of  the  elements  with  which  we 
come  in  contact  on  the  earth ;  the  spectra  of  the  sun  and  the  fixed 
stars  are,  therefore,  absorption  spectra,  caused  by  the  white  light 
of  the  glowing  central  mass  passing  through  the  surrounding  chro- 
mosphere. The  gaseous  envelope  of  the  sun  must  therefore  contain 
the  glowing  vapors  of  a  number  of  elements ;  and  those  elements 
are  identical  with  the  ones  encountered  on  the  earth.*  By  means 
of  the  spectroscope  we  have,  therefore,  been  able,  to  a  great  extent, 
to  analyze  the  composition  of  the  gases  which  surround  the  sun  and 
the  fixed  stars. 

*  Some  lines  appear  in  the  spectrum  of  the  sun  which  do  not  correspond 
to  those  emitted  by  any  known  element  on  the  earth. 


396  COPPER,  SILVER,    AND   GOLD. 


CHAPTEK   LIIL 

COPPER,    SILVER,    AND    GOLD. 

Copper ;  symbol,  Cu;  atomic  weight,  63.6; 
Silver;  symbol,  Ag;  atomic  weight,  107.92 ; 
Gold  ;  symbol,  Au ;  atomic  weight,  197.3. 

THESE  three  elements  find  their  places  at  the  beginning  of  the 
second  section  of  the  long  periods,  and,  because  the  second  portion 
of  the  latter  shows  many  points  of  resemblance  to  the  first,  we  must 
expect  copper,  silver,  and  gold  to  be  in  some  respects  like  the  alkali 
metals,  of  which  family  they  form  the  secondary  group  (see  page 
371).  Naturally,  as  copper,  silver,  and  gold  are  much  nearer  to  the 
not-metallic  end  of  the  long  periods  than  are  the  alkali  metals,  we 
must  not,  in  the  chemical  behavior  of  the  former,  look  for  metallic 
properties  by  any  means  so  pronounced  as  are  encountered  with  the 
latter ;  this  difference  is  manifested  in  marked  degree  by  the  fact 
that  neither  copper,  silver,  nor  gold  decomposes  water;  the  resem- 
blances between  the  elements  of  this  group  and  the  alkalies  are 
confined  chiefly  to  their  univalence,  by  reason  of  which  each  element 
forms  halogen  derivatives  having  the  formula  MX  (corresponding 
to  those  of  the  alkali  metals),  and  to  the  isomorphism  between  the 
crystalline  form  of  some  of  the  compounds  derived  from  these  ele- 
ments in  their  univalent  condition  and  similar  compounds  of  the 
alkalies  (of  the  latter,  especially  those  of  sodium) ;  on  the  other 
hand,  copper  and  gold  differ  very  widely  from  the  alkalies  by  being 
able  to  form  higher  oxides  and  salts  derived  from  these ;  in  the  case 
of  copper,  this  higher  oxide  has  the  formula  Cu  0  ;  of  gold,  Au2  03 . 
The  oxides,  M2  0,  and  the  hydroxides  and  halides  derived  from 
these  are,  therefore,  the  typical  compounds  belonging  to  the  entire 
family  comprising  the  alkali  metals,  as  well  as  copper,  silver,  and 
gold ;  but  those  members  of  this  family  which  find  their  places 
cut  the  beginning  of  the  secondary  division  of  the  long  periods  are  also 
able  to  form  oxides  in  which  the  valence  of  the  elements  is  more 
than  one.  In  the  formation  of  the  latter  oxides,  the  elements  in 


COPPER  :    OCCURRENCE. 


397 


question  (Cu,  Ag,  Au)  appear  as  connecting  links  between  the  ele- 
ments of  the  eighth  group  (Fe,  Co,  Ni,  etc.)  and  the  second  halves  of 
the  long  periods.  Copper,  silver,  and  gold  find  their  places  as  the 
minima  and  the  beginning  of  the  second  portions  of  the  curves  of 
atomic  volumes  (see  page  365) ;  they  are,  therefore,  malleable,  duc- 
tile, fusible,  electropositive,  and  good  conductors  of  electricity. 

Their    chief    physical    constants    are    given*  in   the   following 
table  :  — 


ATOMIC  WEIGHT. 

SPECIFIC  GBAVITY. 

ATOMIC  VOLUME. 

MELTING  POINT. 

Copper, 

63.6 

8.8 

7.1 

1050° 

Silver, 

107.92 

10.5 

10.1 

950° 

Gold, 

197.3 

19.3 

10.2 

1030° 

All  of  these  elements  are  volatilized  when  heated  in  the  flame 
of  the  oxyhydrogeii  blowpipe.  They  crystallize,  as  do  the  alkali 
metals,  in  forms  belonging  to  the  regular  system. 

The  most  important  mineral  forms  in  which  these  elements  occur 
are  given  in  the  following  table :  — 

COPPER. — As  native  copper,  Lake  Superior  region,  Siberia,  Chile,  Aus- 
tralia; as  chalcocite  (cuprous  sulphide),  Cu2S,  in  Cornwall,  Siberia, 
Saxony,  Western  Montana  ;  as  chalcopyrite  (copper  pyrites),  Cu  Fe  S2 , 
similar  localities  to  chalcocite.;  as  cuprite,  Cu2  O,  in  Lake  Superior 
regions;  as  melaconite,  CuO,  in  Lake  Superior  regions. 

Copper  also  frequently  occurs,  combined  with  arsenic  and  anti- 
mony trisulphides,  in  the  mineral  tetrahedrite  (gray  copper  ore, 
fahl-ore),  which  has  approximately  the  formula,  4  (Cu2,  Ag2,  Fe, 
Zn)  S,  Sb2S3,  or  4  (Cu2,  Fe,  Zn)  S,  As2S3.  As  the  formulae  will 
show,  iron  and  zinc  accompany  copper  in  these  ores,  while  silver  is 
also  frequently  encountered  therein.* 

In  addition  to  the  above  sulphur  compounds,  the  basic  carbon- 
ates of  copper,  malachite,  Cu2  (HO)2C03,  and  lazurite,  Cu3  (H0)2 
(C03)2,  are  important  minerals. 

*  In  this  mineral  two  atoms  of  copper  replace  one  of  zinc  or  of  iron  iso- 
morphously.  The  possibility  of  this  replacement  is  probably  due  to  the  fact 
that  the  size  of  the  molecule  prevents  an  undue  influence  on  the  crystalline 
form  by  one  or  two  atoms;  nevertheless,  it  serves  admirably  to  illustrate  the 
necessity  of  great  caution  in  using  the  laws  of  isomorphism  for  the  purpose  of 
determining  atomic  weights  (see  page  356). 


398  COPPER;  METALLURGY. 

SILVER.  —  As  native  silver,  in  the  United  States,  Mexico,  Peru,  Norway, 
Saxony,  Bohemia,  Siberia,  etc. 

As  argentine,  Ag2  S,  isomorphous  with  chalcocite,  Cu2  S  ;  silver 
sulphide  also  occurs  in  conjunction  with  lead  sulphide  and  in  nu- 
merous minerals  in  which  it  is  combined  with  the  sulphides  of 
antimony,  Sb2  S3 ,  arsenic,  As2  S3 ,  and  iron,  Fe2  S3 . 

GOLD.  —  As  native  gold  in  quartz  veins  in  conjunction  with  iron  pyrites, 
chalcopyrite,  galena,  and  other  sulphides.  Gold  particles  also  occur 
in  the  gravel  or  sand  of  rivers  or  valleys  in  auriferous  regions,  or  on 
the  slopes  of  mountains  or  hills  whose  rocks  contain  in  some  part 
auriferous  veins. 

Compounds  of  gold  are  very  infrequent  as  minerals.  The  tellu- 
ride  of  gold  and  silver,  (Ag,  Au)2  Te,  is  sometimes  found. 

METALLURGY  OF  COPPER.  —  The  copper  ores,  the  most  valuable 
of  which  are  the  oxides  and  sulphides,  are  roasted;  by  this  means 
the  volatile  compounds  of  arsenic  and  antimony  are  removed,  while 
the  sulphides  of  iron,  which  are  present,  are  easily  converted  into  the 
oxide,  the  sulphur  passing  off  as  sulphur  dioxide.  If  the  material 
used  in  the  preparation  of  copper  should  contain  large  quantities  of 
quartz,  the  latter  substance,  in  melting,  will  attack  the  ferric  oxide 
in  order  to  form  the  silicate  of  iron,  which  can  be  run  off  in  the 
form  of  slag ;  if  the  ores  used  do  not  already  contain  silicon  diox- 
ide, the  latter  must  be  added  as  quartz  or  sand.  The  product  ob- 
tained after  the  smelting  is  much  richer  in  copper  than  the  original 
ore,  and  contains  both  cuprous  and  cupric  oxides  as  well  as  the 
corresponding  sulphides.  This  material  is  again  roasted  and  the 
remaining  iron  separated  as  slag,  while  the  oxides  and  sulphides  of 
copper  mutually  reduce  each  other  as  follows  :  — 

Cu2  S  +  2  Cua  0  ==  6  Cu  +  S02, 

Cu2S+2CuO  =  4Cu-fS02. 

It  is  frequently  necessary  to  repeat  the  roasting  process  several 
times  while  adding  the  coal  and  sand.  Finally,  the  fused  copper  is 
stirred  with  poles  of  green  wood.  Not  infrequently  the  copper 
ores  contain  considerable  quantities  of  silver,  so  that  the  separation 
of  the  latter  becomes  profitable.  The  processes  used  are,  however, 
somewhat  complicated  and  need  not  be  entered  into  in  this  work. 

SILVER.  —  Considerable  quantities  of  silver  are  obtained  from 
galena  (PbS,  lead  sulphide),  and  separated  from  the  lead  by  the 
process  of  cupellation,  a  description  of  which  was  given  on  pages 


SILVER;  METALLURGY.  399 

320  and  321.  The  relative  amounts  of  silver  in  the  lead  ore  may 
vary  to  such  an  extent  that,  at  one  time,  the  silver  is  the  chief 
product  of  the  process  while,  at  another,  it  is  present  in  such  small 
quantities  that  lead  forms  the  only  commercially  valuable  substance. 
In  some  cases  the  ores  are  ground ;  the  richer  ones  are  roasted  with 
common  salt  so  as  to  convert  the  silver  compounds  into  silver  chlo- 
ride ;  the  crushed  mixture  is  then  treated  with  mercury  and  hot 
water,  either  in  barrels  or  in  cast-iron  pans,  the  mercury  taking  the 
place  of  silver  in  the  chloride  :  — 

AgCl+Hg  =  Ag  +  HgCl, 

while  the  excess  of  the  fluid  metal  dissolves  the  silver  to  form  an 
amalgam.  This  amalgam  is  then  washed,  strained,  and  carefully 
distilled  from  cast-iron  retorts ;  the  mercury  passes  off  and  the  re- 
maining silver  is  cast  into  ingots.  Several  processes  of  silver  ex- 
traction by  the  so-called  "  wet  way "  depend  on  the  conversion  of 
the  silver  ores  into  the  chloride  by  means  of  a  solution  of  sodium 
chloride,  the  extraction  of  the  chloride  of  silver  by  means  of  a  solu- 
tion of  sodium  hyposulphite  (thiosulphate),  in  which  substance  sil- 
ver chloride  is  soluble,  and  the  subsequent  precipitation  of  silver 
sulphide  from  this  solution  by  means  of  the  sulphide  of  sodium. 
The  sulphide  of  silver  is  separated  and  roasted,  by  means  of  which 
process  the  sulphur  burns  off  and  metallic  silver  remains.  Chemi- 
cally pure  silver  is  prepared  by  boiling  pure  silver  chloride  with  a 
mixture  of  potassium  hydroxide  and  grape  sugar. 

GOLD.  —  As  gold  almost  exclusively  occurs  in  the  form  of  the 
native  element,  the  process  of  its  extraction  consists  simply  in  a 
separation  of  the  various  impurities  which  accompany  the  metal. 
Alluvial  washing  (or  placer  digging)  is  done  by  placing  the  aurifer- 
ous deposit  found  on  the  banks  of  rivers  or -in  the  valleys  in  shal- 
low pans,  and  then  washing  off  the  lighter  portions,  while  the 
specifically  heavier  gold  remains  behind,  mixed  with  pebbles  and 
stones.  From  the  latter  it  can  be  mechanically  separated.  The 
former  gold  beds,  having  become  to  a  great  extent  exhausted,  the 
process  of  hydraulic  mining  is  now  frequently  resorted  to.  The  sides 
of  the  hills  which  contain  gold-bearing  conglomerate  are  washed  out 
by  means  of  powerful  streams  of  water ;  the  washings  are  conducted 
through  a  channel  containing  a  number  of  sluice  boxes  which  collect 
the  heavier  particles.  Mercury  is  placed  in  each  of  these  boxes 


400  COPPER;    SILVER;    GOLD;    CHEMICAL   ACTION. 

(because  that  metal  is  capable  of  forming  an  amalgam  with  the 
smaller  particles  of  gold)  ;  the  sluice  boxes  are  opened  from  time  to 
time,  and  the  metal  contained  in  them  is  mechanically  separated. 
Where  the  gold  occurs  in  veins  imbedded  in  quartz,  the  material  is 
mined,  crushed,  and  the  gold  extracted  by  means  of  mercury ;  the 
amalgam  so  formed  is  treated  as  is  that  of  silver. 

Copper,  silver,  and  gold  all  have  a  metallic  lustre  and  are 
malleable  and  ductile  ;  neither  silver  nor  gold  will  burn ;  indeed,  the 
oxides  of  these  metals  decompose  when  heated ;  copper,  on  the 
other  hand,  burns  when  heated  in  oxygen  to  a  high  temperature, 
the  compound  formed  being  cupric  oxide.  Chlorine  and  bromine 
attack  all  three  of  the  metals,  forming  the  corresponding  halides, 
while  iodine  attacks  copper  and  silver. 

Nitric  acid  readily  dissolves  either  copper  or  silver ;  forming 
cupric  nitrate,  Cu  (N03)2,  and  silver  nitrate,  AgN03,  respectively; 
the  acid  does  not  attack  gold  (page  207,  b). 

Aqua  regia  (see  page  203)  attacks  either  copper,  silver,  or  gold, 
producing  the  corresponding  chlorides. 

Sulphuric  acid,  when  hot  and  concentrated,  dissolves  either 
copper  or  silver,  producing  cupric  sulphate,  Cu  S04 ,  and  argentic 
sulphate,  Ag2S04;  sulphur  dioxide  is  at  the  same  time  given  off 
(see  page  136). 

Caustic  alkalies,  in  solution,  hot  and  concentrated,  dissolve  gold. 

The  alloys  of  the  three  metals  are  quite  important.  Two  parts 
of  copper  alloyed  with  one  part  of  zinc  form  a  yellow  metal  (  brass  )  ; 
alloys  of  copper  and  tin  are  known  as  bell  metal,  gun  metal,  and 
bronze,  according  to  the  proportions  of  the  ingredients;  ninety 
parts  of  copper  united  with  ten  parts  of  aluminium  is  the  most 
common  form  of  aluminium  bronze.  Commercial  silver  is  always 
alloyed  with  copper,  as  pure  silver  is  too  soft  for  the  ordinary  pur- 
poses of  coinage  and  the  manufacture  of  jewelry ;  the  silver  coins 
in  use  contain  from  7.5  to  10  per  cent  of  copper ;  gold  is  also 
invariably  alloyed  with  copper  or  silver,  the  resulting  alloys  being- 
much  harder  than  pure  gold.* 

*  The  fineness  of  gold  is  measured  in  carats;  the  number  of  carats  used 
in  designating  a  particular  alloy  of  gold  indicate  the  number  of  parts  of  pure 
gold  contained  in  twenty-four  parts  of  alloy;  thus,  18  carat  gold  has  IS  parts 
of  gold  in  24  parts  of  alloy. 


CUPROUS   COMPOUNDS.  401 

In  their  chemical  behavior,  copper,  silver,  and  gold  differ  quite 
markedly,  and  for  this  reason  it  will  be  necessary  to  discuss  the 
chemistry  of  each  metal  separately. 

Copper  forms  two  oxides,*  cuprous  oxide,  Cu2  0,  and  cupric 
oxide,  Cu  0.  Cuprous  oxide  is  found  in  nature  as  the  mineral  cu- 
prite, occurring  in  octahedra  or  in  cubes  ;  in  the  laboratory  the  oxide 
can  be  produced  by  heating  a  mixture  of  copper  and  cupric  oxide  to 
aredheat:-  Cu  +  CuO  =  Cu2O. 

A  hydroxide  corresponding  to  this  oxide  is  unknown,  but  if 
a  solution  of  copper  sulphate,  mixed  with  glucose,  is  warmed  with 
alkalies,  a  yellow  precipitate  which  has  the  formula  4  Cu2  0  -f-  H2  0 
is  produced ;  this  substance  is  not  completely  dehydrated  until  a  tem- 
perature of  360°  is  reached  ;  when  exposed  to  the  air  this  hydrate 
rapidly  takes  up  oxygen,  forming  blue  cupric  hydroxide,  Cu  (OH)2. 
Cuprous  oxide  dissolves  in  ammonia  to  produce  a  colorless  fluid, 
which,  however,  soon  absorbs  oxygen  from  the  atmosphere  and 
assumes  a  deep  blue  color.  Cuprous  oxide  is  very  easily  decom- 
posed by  many  oxy-acids,  and  for  this  reason  very  few  cuprous 
salts  are  known;  dilute  nitric  or  sulphuric  acids  attack  it,  liber- 
ating metallic  copper  and  producing  the  corresponding  cupric 

Cu2  0  +  H2  S04  =  Cu  S04  -f  H2  0  +  Cu. 

The  most  important  cuprous  salts  are  the  cuprous  halides. 

Cuprous  chloride,  Cu  Cl,  is  a  white  solid  which  is  with  difficulty 
soluble  in  water.  In  this  respect  it  resembles  the  corresponding 
chloride  of  silver.  Like  the  latter  it  is  readily  dissolved  by 
ammonia,  with  which  it  forms  a  substance  having  the  formula 
NH3CuCl;  the  latter  compound  possibly  consists  of  ammonium 
chloride,  NH3  H  Cl,  with  the  difference  that  one  atom  of  hydrogen 
in  the  formula  weight  has  been  replaced  by-  one  of  copper,  so  that 
in  the  formation  of  this  body  cuprous  chloride  would  act  as  does 
hydrochloric  acid  in  ammonium  chloride.  Cuprous  chloride  is 
readily  oxidized  when  exposed  to  the  air. 

Cuprous  iodide  is  the  only  iodide  of  copper  which  exists.  It  is 
formed  by  adding  the  solution  of  an  iodide  to  a  copper  sulphate 
solution ;  the  cupric  iodide,  which  we  should  expect  to  be  formed, 
at  once  breaking  down  into  cuprous  iodide  and  iodine  :  — 

*  An  oxide  Cu4  O  has  also  been  described. 


402  CUPRIC   COMPOUNDS. 

Cu  S04  +  2  KI  =  K2  S04  +  Cu  I  +  I.* 

Cupric  oxide,  Cu  0,  is  the  most  stable  oxide  of  copper.  It  can 
readily  be  formed  by  heating  copper  in  a  current  of  air  or  of  oxygen, 
or  by  decomposing  cupric  nitrate,  Cu  (N03)2,  (page  208,  •/?).  It  is 
a  black  substance  which  readily  loses  oxygen  when  it  is  heated  in  a 
current  of  hydrogen  (see  page  39)  or  with  charcoal :  — 

Cu  0  -f  C  =  Cu  -f  CO. 

Cupric  hydroxide,  Cu  (OH)2,  is  formed  in  a  manner  parallel  to 
the  formation  of  the  hydroxides  of  most  metals  i.e.,  by  precipitation 
from  the  solution  of  a  copper  salt  upon  the  addition  of  a  soluble 
hydroxide.  It  appears  as  a  blue,  flaky  precipitate,  which  readily 
loses  water  when  it  is  warmed  ;  it  then  turns  black  and  forms  cupric 
oxide.  Ammonia  water  dissolves  both  the  oxide  and  hydroxide; 
the  solution  has  a  deep  blue  color  which  can  be  observed  even  when 
only  very  little  copper  is  present.  Both  cupric  oxide  and  hydroxide 
are  bases ;  they  dissolve  in  acids  to  form  stable  cupric  salts.  The 
latter,  when  they  contain  water  of  crystallization,  are  generally  blue 
or  green. 

Cupric  chloride,  Cu  C12 ,  is  formed  by  the  action  of  chlorine  on 
copper,  or  by  dissolving  the  oxide  or  hydroxide  in  hydrochloric 
acid,  and  then  evaporating  the  solution  and  drying  at  100° ;  when 
anhydrous,  cupric  chloride  is  a  brown  powder ;  when  crystallized 
from  water  it  forms  green  crystals  of  the  composition  Cu  C12  -f-2  H2  0 ; 
it  is  readily  soluble  in  water.  The  chloride  forms  double  salts  with 
the  chlorides  of  the  alkali  metals.  The  bromide  resembles  the 
chloride  in  every  respect.  , 

Cupric  sulphate,  Cu  S04  +  5  H2  0  ( blue  vitriol),  is  formed  by 
dissolving  copper  in  sulphuric  acid  (page  137). t  The  commercial 
product  is  prepared  by  heating  copper  sulphide  in  a  current  of  air, 
extracting  with  water  and  recrystallizing.  It  forms  large,  blue 
crystals  belonging  to  the  triclinic  system ;  it  is  readily  soluble  in 
water.  When  exposed  to  the  air  the  crystals  effloresce  and  lose 
two  molecules  of  water  of  crystallization ;  at  100°  two  more  mole- 

*  The  iodine  which  is  liberated  can  be  removed  by  the  addition  of  a 
reducing  agent  such  as  SO2  (see  page  139). 

+  In  this  reaction  several  secondary  products  [cuprous  sulphide,  Cu.2S, 
and  compounds,  (CuaS,  CuO,)  and  (CuS,  CuO)]  are  produced. 


CUPRIC    SALTS.  403 

cules  pass  off,  so  that  a  salt  of  the  composition  Cu  S04  +  H2  0  is 
left ;  the  latter  is  probably  a  secondary  salt  of  the  hydrated  sul- 
phuric acid  having  the  formula  H4  S05 ,  and  should  therefore  be 
written  CuH2S05.  At  230°  the  two  hydroxyl  groups  present  in 
this  salt  (see  page  153)  finally  separate  water,  and  leave  a  white 
powder  which  has  a  composition  expressed  by  the  formula  Cu  SO4 ; 
as  soon  as  water  is  added  to  this,  a  blue  solution  containing 
Cu  S04-|-5H2O  is  produced;  the  latter  belongs  to  a  class  of  sul- 
phates known  as  the  vitriols  (see  magnesium).  Copper  sulphate  can 
unite  with  ammonia  to  form  compounds  in  which  molecules  of 
ammonia  take  the  place  of  molecules  of  water  of  crystallization,* 
for  instance,  salts  having  the  composition  Cu  S04  +  5  NH3 , 
Cu  S04  +  4  ISTHg  +  H2  0,  Cu  S04  +  3  NH3  +  2  H2  0,  etc.,  are  capa- 
ble of  existence.  These  substances  furnish  examples  of  cases  where 
the  compound  of  nitrogen  and  hydrogen  plays  the  same  role  as  the 
compound  of  oxygen  and  hydrogen  (see  page  299).  An  adequate 
chemical  explanation  of  the  existence  of  salts  with  water  of  crys- 
tallization and  of  these  ammonium  compounds,  provided  we  adhere 
to  our  present  theories  of  valence,  is  as  yet  lacking. f  A  number  of 
basic  sulphates  of  copper  have  been  described. 

Cupric  nitrate,  Cu  (N03)2 -f  3  H2  0,  is  formed  by  dissolving 
copper  in  nitric  acid  (see  page  198)  or  by  adding  nitric  acid  to 
cupric  oxide  or  hydroxide.  The  salt  exists  in  blue,  prismatic  crys- 
tals which  are  soluble  in  water,  and  which  break  down  completely 
into  cupric  oxide  and  nitrogen  peroxide  when  heated  (see  page  208). 

Basic  carbonates  of  copper.  The  secondary,  normal  carbonate, 
Cu  C03 ,  is  unknown.  A  basic  carbonate  which  has  the  same  com- 
position as  the  mineral  lazurite  is  produced  by  adding  the  solution 
of  the  carbonate  of  an  alkali  metal  to  a  solution  of  a  copper  salt. 

*  The  same  is  true  of  cupric  chloride,  Cu  CL>  +  2  H2  O,  for  a  substance  of 
the  composition  Cu  C12  +  2  NH3  is  known. 

t  The  nitrogen  atom  in  ammonia  is  unsaturated,  as  is  proved  by  the  easy 
formation  of  ammonium  compounds.  The  resemblance  between  ammonia  and 
water  in  these  salts  would  lead  us  to  suspect  that  the  oxygen  atom  is  also  un- 
saturated in  water.  This  conclusion  is  not  very  startling  if  we  take  into  con- 
sideration the  great  resemblance  between  oxygen  and  sulphur,  the  latter 
element  being  able  to  take  part  in  compounds  in  which  its  atoms  are  hexa- 
valent.  It  has  also  been  shown  that  methyl  ether  (CH3  —  O  —  CH3)  can  add 
hydrochloric  acid,  just  as  ammonia  would;  the  unstable  compound  so  formed 
probably  contains  tetravalent  oxygen. 


404  COPPER;    SULPHIDES. 

The  carbonate  is  insoluble  in  water,  and  has  the  composition 
Cu(OH)2CuC03or:-  Cu_OH 

)C03 
Cu  — OH. 

Other  basic  carbonates  (malachite,  for  example)  occur  in  the  form 
of  minerals.  All  copper  carbonates,  when  heated  to  300°,  lose  car- 
bon dioxide  and  leave  cupric  oxide.  The  green  coating  formed 
on  metallic  copper  which  has  been  exposed  to  moist  air,  is  due  to 
the  formation  of  a  basic  carbonate. 

Sulphides  of  copper.  Copper  forms  two  sulphides  which,  in 
formula,  correspond  to  the  oxides.  Cuprous  sulphide,  Cu2  S,  occurs 
as  the  mineral  chalcocite  and  is  also,  most  probably,  a  constituent  of 
chalcopyrite,  for  the  latter  substance  is  regarded  as  a  sulpho-salt  in 
which  cuprous  sulphide  is  the  base  and  ferric  sulphide  the  acidic 
anhydride:-  C^  +  ^  =  Gn^t* 

Cuprous  sulphide  is  a  substance  which  is  formed  with  remarkable 
readiness,  for  it  can  be  produced  by  merely  pressing  copper  and  sul- 
phur firmly  together,!  or  by  heating  or  even  rubbing  copper  with  a 
sufficient  quantity  of  sulphur  to  form  cuprous  sulphide. 

*  Cupric  sulphide.!  Cu  S,  is  produced  by  precipitating  a  slightly 
acid  solution  of  a  copper  salt  by  means  of  hydrogen  sulphide  (see 
page  100).  Cupric  sulphide  dissolves  in  hot,  concentrated  hydrochlo- 
ric acid,  forming  hydrogen  sulphide  and  cupric  chloride.  Nitric 
acid  dissolves  it,  producing  cupric  nitrate,  Cu  (N03)2,  while  the 
hydrogen  sulphide,  which  is  liberated,  is  oxidized  to  sulphur  (see 
page  206).  The  precipitate  is  nearly  black,  and,  when  moist,  is 
readily  oxidized  to  cupric  sulphate  when  it  is  exposed  to  the  air :  —  $ 

Cu  S  +  4  0  =  Cu  S04 . 

When  heated  to  200°,  in  a  current  of  hydrogen,  it  is  changed  into 
cuprous  sulphide. 

*  One-half  the  formula  weight  of  the  above  salt  would  lead  to  the  usual 
formula  of  chalcopyrite,  namely,  Cu  Fe  S2  . 

t  This  sulphurizing  action  reminds  us  forcibly  of  the  oxidation  of  copper 
in  the  air. 

f  This  oxidation  is  taken  advantage  of  in  the  commercial  formation  of 
blue  vitriol  from  sulphides  of  copper. 


SILVER  ;    SALTS    OF.  405 

Compounds  of  silver.  Silver  is  univalent  in  nearly  all  of  its 
compounds.  Its  most  stable  oxide  is  Ag2  0,  but  oxides  having  the 
formulae  Ag4  0  *  and  Ag  0  are  also  known.  The  most  stable  oxide, 
Ag2  0,  is  produced  by  adding  a  soluble-  hydroxide  to  a  solution  of  a 
silver  salt.  It  is  a  black  precipitate  which  readily  dissolves  in  an 
excess  of  ammonia  (compare  with  cuprous  oxide).  When  heated  it 
readily  breaks  down  into  silver  and  oxygen. 

Silver  chloride,  Ag  Cl,  one  of  the  most  characteristic  silver  salts, 
is  insoluble  in  water  and  is  therefore  produced  as  a  white  precipitate 
when  hydrochloric  acid,  or  the  solution  of  a  chloride,  is  added  to  a 
solution  of  a  silver  salt  :  — 


AgN03  +  H    Cl    =  HN03    +AgCl,t 
Ag  N03  +  Na  Cl    =  .Na  N03  +  Ag  Cl. 
Soluble.         Soluble.  Soluble.         Insoluble. 

Owing  to  this  insolubility,  the  formation  of  silver  chloride  affords 
a  ready  means  of  detecting  the  presence  either  of  silver  or  of  a 
chloride  in  a  solution.  When  silver  chloride  is  exposed  to  the  light, 
it  rapidly  changes  color,  becoming  violet  at  first  ;  that  color  however 
soon  becomes  darker  until  the  entire  mass  turns  black.  The  same 
change  also  takes  place  with  the  equally  insoluble  bromide  or  iodide 
of  silver,  which  latter  salts  can  be  formed  by  precipitation  exactly 
as  is  the  chloride.  The  chloride  of  silver  is  sometimes  found  as  a 
mineral,  the  name  of  which  is  cerargyrite.  Silver  chloride  is  very 
readily  dissolved  by  an  aqueous  solution  of  ammonia  ;  silver  bromide 
is  less  soluble  in  that  medium,  while  silver  iodide  is  entirely  insolu- 
ble. When  heated  with  hydrochloric  acid,  both  the  bromide  and 
iodide  of  silver  are  converted  into  the  chloride,  but,  on  the  other 
hand,  the  chloride  of  silver  changes  into  the  bromide  when  treated 
with  cold  hydrobromic  acid,  and  both  the  chloride  and  bromide  are 
converted  into  the  iodide  when  covered  with  cold  hydroiodic  acid 
and  allowed  to  stand.  $ 

One  of  the  chief  uses  for  the  silver  halides  is  in  the  art  of  pho- 
tography ;  their  application  depends  on  the  fact  that  light  effects 

*  The  existence  of  this  oxide  is  doubtful  ;  it  is  possibly  a  mixture  of  metal- 
lic silver  and  of  Ag2  O. 

t  See  page  380. 

t  These  changes  afford  an  excellent  example  of  the  reversability  of  many 
chemical  reactions  when  the  conditions  are  changed. 


406  PHOTO-CHEMICAL   ACTION. 

such  a  change  in  these  compounds  as  to  cause  the  subsequent  for- 
mation of  a  film  of  finely  divided  metallic  silver  upon  those  portions 
of  a  glass  plate  which  have  been  covered  with  a  thin  layer  of  silver 
halide  and  which  have  been  acted  upon  by  the  light ;  this  film  being 
produced  when  the  silver  halides  are  treated  with  certain  reducing 
solutions ;  *  the  unchanged  silver  halide  is  subsequently  removed 
by  solutions  which  have  a  solvent  action  on  the  unchanged  salts  but 
which  leave  metallic  silver  untouched.  A  solution  of  this  nature 
is,  for  instance,  one  of  sodium  thiosulphate  (hyposulphite),  Na2  S2  Oa 
(see  page  154).  | 

A  plate  covered  with  a  thin  layer  of  gelatine,  containing  some  sil- 
ver halide  (preferably  the  iodide)  very  evenly  distributed  through- 
out the  mass,  is  exposed  to  the  light  in  such  a  manner  that  a 
perfectly  clear  image  of  some  object  to  be  photographed  is  thrown 
upon  it  by  means  of  a  camera.  The  changes  in  the  silver  salt  then 
take  place  according  to  the  intensity  of  the  light  which  is  cast  upon 
the  object.  When  the  silver  is  finally  deposited  on  the  plate  by 
means  of  the  processes  which  have  been  outlined  above,  a  perfectly 
clear  image  of  the  object  to  be  photographed  is  left.  The  plate  so 
produced  is  a  "  negative ; "  the  photograph  is  prepared  from  this  by 
covering  paper  with  a  sensitive  film  similar  to  that  which  was  used 
on  the  negative,  and  then  exposing  a  piece  of  this,  placed  under  the 
negative  upon  which  the  image  has  been  fixed,  to  the  action  of  the 
light ;  the  picture  is  then  produced  in  such  a  manner  that,  wherever 
a  dark  spot  appears  on  the  negative,  a  light  spot  will  be  seen  on  the 
photograph.  The  latter  is  subsequently  developed  and  fixed  by  a 
process  similar  to  that  used  in  the  preparation  of  the  negative. 

The  chemical  action  of  light  on  the  silver  halides  is  not  as  yet 
definitely  understood.  The  chemical  changes  which  are  useful  in 
the  art  of  photography  are  not,  however,  by  any  means  the  only 
ones  which  are  caused  by  light  and  which  we  encounter.  Hydro- 
gen and  chlorine,  it  will  be  remembered,  are  entirely  without  action 
on  each  other  when  in  the  dark,  but  unite  with  the  greatest  readi- 
ness in  the  sunlight ;  the  compounds  of  carbon  and  hydrogen  (see 
page  277)  are  substituted  by  halogens  when  placed  in  the  light, 

*  So-called  developers  —  ferrous  sulphate,  pyrogallic  acid,  liydrochinon,  etc. 
t  The  cause  of  the  solution  is  the  formation  of  the  soluble  double  thiosul- 
phate of  silver  and  sodium,  Ag2  S2  O3  ,  Na2  S2  O3 . 


SILVER  ;  NITRATE,  SULPHIDE.  407 

but  are  not  attacked  in  the  dark,  while  many  vegetable  dyes  bleach 
wii^k  the  greatest  rapidity  when  under  the  influence  of  the  sunlight. 
The  action  of  light  in  bringing  about  chemical  changes  is  identical 
with  that  of  heat ;  and,  probably,  all  light  rays  are  capable  of  causing 
such  changes.  It  was  formerly  supposed  that  only  certain  light  rays 
were  capable  of  causing  chemical  reactions  and  these  were  desig- 
nated as  "actinic"  rays;  in  view  of  later  developments  the  distinc- 
tion between  "  actinic "  and  "  not-actinic "  rays  has,  of  course, 
disappeared.  4 

Silver  nitrate,  Ag  N03 ,  is  formed  by  dissolving  silver,  or  silver 
oxide,  in  nitric  acid ;  it  is  soluble  in  water,  and  is  separated  from 
its  solution  by  evaporation.*  The  salt  crystallizes  in  plates  belong- 
ing to  the  rhombic  system,  and  can  be  fused  at  224°  without  chan- 
ging its  composition ;  the  fused  salt  is  cast  into  sticks,  and  is  popu- 
larly termed  "  lunar  caustic."  Silver  nitrate  is  not  changed  when 
exposed  to  the  light,  unless  some  organic  substances  (dust,  etc.)  are 
brought  in  contact  with  it ;  it  then  turns  black,  owing  to  reduction 
and  deposition  of  metallic  silver.f  Dry  silver  nitrate  absorbs 
ammonia  quite  readily ;  the  same  is  true  of  the  dry  chloride.  The 
latter  substance  forms  a  compound,  2  Ag  Cl  +  3  NH3 ,  $  which  loses 
ammonia  when  heated.  This  power  of  absorbing  ammonia  is  one 
of  the  most  marked  resemblances  between  silver  and  copper  salts. 

Sulphide  of  silver,  Ag2  S,  is  formed  by  fusing  sulphur  and  silver 
together,  or  by  precipitation  from  the  solution  of  a  soluble  silver 
salt  by  means  of  hydrogen  sulphide ;  the  natural  mineral  argentite, 
Ag2  S,  has  a  metallic  appearance  much  like  that  of  lead ;  the  pre- 
cipitated sulphide  is  black.  The  compound  is  readily  decomposed 
by  heating  with  lead  or  iron ;  by  this  process  the  sulphides  of  the 
latter  metals  and  free  silver  are  produced.  The  blackening  of  silver 
when  exposed  to  the  air  is  brought  about  by  the  action  of  hydrogen 
sulphide ;  the  minute  traces  of  that  gas  which  are  present  in  the 
atmosphere  attack  the  metal  and  form  silver  sulphide. 

Many  other  salts  of  silver  have  been  investigated.     A  number  of 

*  Silver  nitrate  in  a  pure  form  can  be  precipitated  from  its  concentrated 
aqueous  solutions  by  the  addition  of  concentrated  nitric  acid,  in  which  sub- 
stance it  is  insoluble,  or  nearly  insoluble. 

t  Silver  nitrate,  when  in  contact  with  the  skin,  produces  a  black  stain. 

t  Other  compounds,  AgCl  +  3NH3  and  Ag  Cl  +  2  NH3 ,  have  also  been 
described. 


408  GOLD  ;    SALTS    OF. 

these  display  a  tendency  to  form  double  salts  with  other  silver  com- 
pounds, or  with  the  salts  of  the  alkali  metals.  For  their  descrip- 
tion a  larger  work  must  be  consulted. 

Compounds  of  gold.  Gold  forms  three  oxides,  aurous  oxide, 
Au2  0,  aurous-auric  oxide,  Au4  04 ,  and  auric  oxide,  Au2  Os .  The 
first  of  these  is  produced  with  great  difficulty,  and  is  of  no  great 
importance  excepting  as  an  illustration  of  the  resemblance  between 
the  compounds  of  gold,  silver,  and  copper.  The  last  one,  Au2  03  is 
both  a  base  and  an  acid.  The  hydroxide,  Au(OH)3,  loses  water 
when  allowed  to  stand,  changing  to  a  metahydroxide  of  the  formula 
AuO(OH);  the  latter  changes  to  auric  oxide  at  150°;  at  220° 
auric  oxide  decomposes  completely  into  gold  and  oxygen.  The 
greatest  difficulty  which  is  encountered  in  the  study  of  gold  com- 
pounds lies  in  the  readiness  with  which  they  break  down  and  sepa- 
rate metallic  gold.  Auric  oxide  is  readily  dissolved  by  potassium 
or  sodium  hydroxide ;  the  salts  produced  are  derived  from  meta- 
auric  hydroxide,  Au  02  H,  so  that  they  have  the  formula,  M  Au  02 ; 
in  this  respect  auric  oxide  resembles  the  oxide  of  aluminium  (see 
page  338). 

Chlorides  of  gold.  Three  chlorides  of  gold,  corresponding  to 
the  oxides,  are  known.  They  are  AuCl,  aurous  chloride;  Au2Cl4, 
aurous-auric  chloride ;  and  Au  C18,  auric  chloride.  The  first  of 
these,  Au  Cl,  is  insoluble  in  water.  When  the  substance  is  covered 
with  water  and  allowed  to  stand,  it  breaks  down  into  auric  chloride 
(which  is  soluble)  and  metallic  gold  :  — 

3  Au  Cl  =  Au  C13  +  2  Au ; 

but,  on  the  other  hand,  a  solution  of  auric  chloride,  when  evapo- 
rated, breaks  down  into  aurous  chloride  and  free  chlorine. 

Auric  chloride.  Gold  is  dissolved  by  aqua  regia ;  the  substance 
contained  in  solution  is  auric  chloride,  but  the  latter  cannot  be 
isolated  by  evaporation,  because,  as  was  just  mentioned,  it  decom- 
poses into  aurous  chloride  and  chlorine.  Gold,  when  finely  divided 
and  treated  with  dry  chlorine  at  a  temperature  of  180°,  forms  aurous- 
auric  chloride,  Au2  C14 ;  this  salt,  when  heated  to  220°,  breaks 
down  into  gold  and  auric  chloride  ;  the  auric  chloride  sublimes  and 
collects  on  the  cooler  surfaces.  Chlorine  does  not  attack  gold  at 
300°.  A  solution  of  auric  chloride  can  also  be  prepared  by  allow- 
ing aurous  chloride  (covered  with  water)  to  stand  (see  above)  ; 


CHLOKAURIC   ACIDS. 


409 


when  this  solution  is  carefully  evaporated,   crystals    having  the 
composition  Au  C13  -f-  2  H2  0  are  formed. 

Chlorauric  acids.  When  a  solution  of  gold  in  aqua  regia,  to 
which  concentrated  hydrochloric  acid  has  been  added,  is  evaporated, 
crystals  of  an  acid,  H  Au  C14  -j-  4  H2  O,  separate.  This  compound 
is  another  one  of  the  class  of  substances  of  which  fluosilicic  and 
fluoboric  acids  (pages  303,  330)  and  the  double  salts  of  aluminium 
are  examples  (see  page  337),  two  chlorine  atoms  taking  the  place 
of  one  oxygen  atom ;  the  parallelism  becomes  clear  when  we  com- 
pare the  formula  of  potassium  aurate  with  those  of  chlorauric  acid 
and  of  its  potassium  salt :  — 

0  .     fCl,  ...((31, 


AUJOK  Aujci*H         Au(^ 

Potassium  aurate.        Chlorauric  acid.        Potassium  chloraurate. 

Potassium  chloraurate,  KAuCl4,  can  be  produced  by  mixing 
solutions  of  auric  chloride  and  potassium  chloride  and  evaporating 
to  crystallization.  Potassium  bromoaurate,  KAuBr4,  and  potas- 
sium iodoaurate,  K  Au  I4 ,  are  also  known. 

Two  sulphides  of  gold,  aurous  sulphide,  Au2  S,  and  auric  sul- 
phide, Au2  S3 ,  are  known.  The  former  is  produced  by  treating  a  so- 
lution of  potassium  aurous  cyanide  (Au  C  N,  KCN  =  K  Au  (CN  )2) 
with  hydrogen  sulphide ;  the  latter  by  precipitation  from  a  cold 
neutral  solution  of  auric  chloride  by  means  of  hydrogen  sulphide ; 
if  the  solution  is  hot,  nothing  but  metallic  gold  separates.* 

The  chief  compounds  discussed  in  the  last  chapter  are  given  in 
the  following  tables  :  — 

COMPOUNDS     TYPICAL    OF     THE    FAMILY,    CORRESPONDING     TO    COM- 
POUNDS   OF    THE    ALKALI    METALS. 


OXIDES. 

CHLORIDES. 

6DLPHIDE8. 

Copper, 

Cu2O 

CuCl  t 

Cu2Sf  - 

Cuprous  compounds. 

Silver, 

A&O: 

AgClf 

Ag2Sf 

Argentic  compounds. 

Gold, 

Au2OJ 

AuClf 

Au2Sf 

Aurous  compounds. 

The  compounds  of  copper  and  silver  unite  with  ammonia. 
The  salts  of  silver  are  derived  from  the  oxide  Ag3  O ;  for  instance,  Ag  NO3 , 
Ag3S04. 

t  Insoluble  in  water  and  in  dilute  acids. 

t  Decomposed  into  oxygen  and  the  metal  when  heated. 


A  gold  sulphide,  Au  S,  is  also  described. 


410  COPPER  ;    SILVER  ;   GOLD  ;   COMPOUNDS   OF. 

COMPOUNDS    NOT    TYPICAL. 


OXIDES. 

HYUBOXIDES. 

CHI.OBIDE8. 

SULPHIDES. 

Cupric  compounds, 
Auric  compounds, 

CuO 
Au203t 

Cu(OH)2 

Au(OH)3t 

CuCl.2* 
AuCl3*t 

CuS 

Au2  S3  t 

The  salts  of  copper  are  derived  from  cupric  oxide;  for  instance  Cu  (NO3)2 , 
Cu  SO4 . 

Auric  oxide  is  acidic  in  its  character.  It  is  also  basic,  for  the  few  gold 
salts  which  are  known  (except  AuCl)  are  derived  from  it.  The  aurates  are 
derived  from  the  hydroxide  Au  O2  H ;  the  salts  derived  from  the  alkalies  are 
Au  O2  M.  Auric  chloride  has  the  character  of  an  acidic  anhydride ;  chlorau- 
ric  acid,  Au  CU  H,  chloraurates,  Au  CU  M. 

*  Soluble  in  water. 

t  Decomposed  by  heat,  leaving  gold  behind. 


'if  i 


THE   ALKALINE   EARTHS.  411 


CHAPTER   LIV. 

THE  FAMILY  OF  THE  ALKALINE  EARTHS. 

Beryllium  (Glucinum)  ;  symbol,  Be  ;   atomic  weight,  9  ; 
Magnesium  ;  symbol,  Mg  ;  atomic  weight,  24.3  ; 
Calcium  ;  symbol,  Ca  ;  atomic  weight,  40  ; 
Strontium  ;  symbol,  Sr  ;  atomic  weight,  87.6  ; 
Barium  ;  symbol,  Ba  ;  atomic  weight,  137.43. 

THE  elements  comprising  the  primary  group  of  the  elements  of 
the  family  of  alkaline  earths  are  :  — 

Beryllium,   atomic  weight,   9  specific  gravity,  1.99,  atomic  volume,  4.52 

Magnesium,     "           "        24.3  "             "        1.74        "            "        13.8 

Calcium,           "           "        40  "             "        1.57        "            "         25.4 

Strontium,       "           "        87.6  "             "        2.50       "            "         34.9 

Barium,           "           "      137.43  "             "        3.75       "           "        36.5 


Of  these,  the  first  two  (beryllium  and  magnesium)  belong  to  the 
typical  short  periods,  and  therefore  resemble  the  following  three 
(calcium,  barium,  and  strontium),  but  they  also  are  closely  allied 
with  the  three  elements  (zinc,  cadmium,  and  mercury)  which  com- 
prise the  secondary  group  belonging  to  this  family.  They  resemble 
calcium,  barium,  and  strontium,  because  their  oxides,  hydroxides, 
and  the  salts  derived  from  them  are  formed  according  to  the  same 
formulae;  they  differ  from  those  three  elements  and  fall  into  line 
with  zinc,  cadmium,  and  mercury,  by  reason  of  the  solubility  of 
their  sulphates  *  and  because  of  their  tendency  to  form  double 
salts  when  their  salts  are  brought  in  contact  with  those  of  ammo- 
nium. 

Beryllium  f  and  magnesium  are  prepared  as  metals  by  heating 
the  chlorides  with  metallic  sodium  :  — 

*  The  sulphates  of  calcium,  strontium,  and  barium,  are  insoluble  or 
nearly  insoluble  in  water. 

t  Beryllium  chiefly  occurs  as  a  metasilicate  of  beryllium  and  aluminium, 
known  as  beryl,  Be3  A12  (Si  O3)6.  When  these  crystals  are  transparent  and 
colored  green  by  chromic  oxide  they  are  termed  emerald. 


412  ALKALINE   EARTHS;    PROPERTIES    OF. 

MCl2+2Na==2NaCl  +  M. 

This  method  is  exactly  parallel  to  the  one  formerly  employed  in 
the  preparation  of  aluminium  (see  page  333).  The  name  glucinum 
was  first  given  to  beryllium,  owing  to  the  sweetish  taste  of  the 
salts  of  this  metal. 

Beryllium  is  white  with  a  silver-like  lustre,  malleable  and  ductile;  its 
melting  point  is  lower  than  that  of  silver;  it  is  slightly  oxidized  when  heated 
to  a  high  temperature ;  when  finely  powdered  and  heated  it  Tmrns  with  a  bril- 
liant light;  it  readily  burns  in  an  atmosphere  of  chlorine;  it  dissolves  in  aque- 
ous hydrochloric  acid. 

Magnesium  is  silver  white,  malleable  and  ductile,  melts  at  700°,  is  volatile 
.at  a  high  heat;  the  metal  does  not  oxidize  in  dry  air,  but  it  readily  corrodes 
when  in  contact  with  water;  when  heated  above  its  melting  point  in  the  air  it 
burns  with  a  most  brilliant  white  light;*  the  metal  also  readily  burns  in 
chlorine;  it  is  easily  soluble  in  dilute  acids. 

Calcium,  barium,  and  strontium  are  isolated  by  electrolyzing  the 
fused  chlorides  in  a  crucible  from  which  air  is  excluded ;  the  posi- 
tive electrode  is  made  of  gas  carbon,  which  is  not  attacked  by 
chlorine,  the  negative  electrode  consists  of  iron  wire.  This  method 
is  also  employed  in  preparing  the  alkali  metals.  Calcium  can  also 
be  made  by  reducing  the  iodide  by  means  of  metallic  sodium. 

Calcium,  strontium,  and  barium  are  yellow  with  metallic  lustre.  The 
freshly  cut  surfaces  of  the  metals  soon  become  covered  with  a  layer  of  oxide ; 
the  metals  must  therefore  be  preserved  under  petroleum  (see  page  384). 
When  heated  in  the  air  they  burn  with  a  brilliant  light.  They  all  energeti- 
cally decompose  water  at  ordinary  temperatures,  liberating  hydrogen. 

The  changes  which  are  brought  about  by  the  increasing  atomic 
weights  and  volumes,  as  we  pass  downward  in  the  list  given  above, 
are  shown  in  the  following  table  :  — 

Beryllium  does  not  decompose  water. 

Magnesium  decomposes  boiling  water. 

Calcium  decomposes  water  at  ordinary  temperatures. 

Strontium  decomposes  water  very  readily  at  ordinary  temperatures. 

Barium  decomposes  water  as  readily  as  the  alkali  metals. 

The  oxides  and  hydroxides  which  are  typical  of  the  family  have 
the  formulae  MO  and  M(OH)2;  the  metals  are  therefore  divalent 
and  replace  two  atoms  of  hydrogen  in  acids. 

*  This  magnesium  light  has  a  marked  chemical  effect  on  silver  halides, 
and  is  used  as  a  flash  light  in  photography. 


CALCIUM;  OXIDE,  HYDROXIDE. 


413 


OXIDES. 

HYDROXIDES. 

Beryllium 

BeO 

Not  changed  to  hydroxide  by  the  ad- 
dition of  water 

Be  (OHa) 

Magnesium 

MgO 

Slowly  changed  to  hydroxide  by  the 
addition  of  water 

Mg(OH)2 

Calcium 

CaO 

Readily  changed  to  hydroxide  by  the 
addition  of  water 

Ca(OH)2 

Strontium 

SrO 

Readily  changed  to  hydroxide  by  the 
addition  of  water 

Sr(OH)2 

Barium 

BaO 

Readily  changed  to  hydroxide  by  the 
addition  of  water 

Ba(OH)2 

The  hydroxides  of  beryllium  and  of  magnesium  are  insoluble  in 
water.  The  solubility  of  the  other  three  increases  with  increasing 
atomic  weight  (see  page  385)  while  the  stability  increases  with  the 
solubility. 

Be  (OH)2  decomposes  at  about  300°;          Be  (OH)2    =  Be  O  +  H2  O. 

Mg  (OH)a        "  "  a  low  red  heat;   Mg  (OH)2  =  Mg  O  +  H.2  O. 

Ca(OH)2         "  "  a  highredheat;  Ca(OH)2  =  Ca  O  +H2O. 

Sr(OH)2          "  "a  white  heat;       Sr(OH)2    =  SrO+H2O. 

Ba  (OH)2,  which  crystallizes  from  water  in  crystals  of  the  formula 
Ba  (OH)2  +  8  H2O,  and  which  is  quite  soluble  in  water,  can  be  fused  without 
change. 

The  oxides  and  hydroxides  are  all  strong  bases;  the  solutions  of  calcium, 
strontium,  and  barium  hydroxides  have  an  alkaline  reaction.  The  hydroxide 
of  beryllium,  being  that  of  the  least  metallic  of  all  the  elements,  can  also  dis- 
solve in  caustic  alkalies,  so  that,  under  certain  circumstances,  it  acts  as  an  acid. 
The  oxide  and  hydroxMe  of  calcium,  known  respectively  as  quick  and  slaked 
lime,  are  the  most  important  of  these  compounds. 

Calcium  oxide  is  prepared  by  heating  the  carbonate  in  "  lime- 
kilns "  until  it  decomposes  into  calcium  oxide  and  carbon  di- 

Ca  C03  =  Ca  0  +  C02 . 

The  quick-lime  so  prepared  is  more  or  less  impure,  according  to  the 
condition  of  the  limestone  or  marble  used ;  when  brought  in  contact 
with  water  it  unites  with  that  liquid  to  form  calcium  hydroxide 
(or  slaked  lime) ;  at  the  same  time  a  large  amount  of  heat  is 

developed :  — 

CaO+H20  =  Ca(OH)2. 

If  quick-lime  is  exposed  to  the  air,  it  absorbs  carbon  dioxide  and 
water.  It  then  crumbles  and  is  said  to  be  "  air  slaked." 


414  BARIUM   SUPEROXIDE. 

Slaked  lime  finds  its  chief  use  in  the  preparation  of  mortar. 
Mortar  is  prepared  by  stirring  together  slaked  lime  and  sand  until 
the  mass  assumes  the  consistency  of  thick  porridge.  When  placed 
between  bricks  the  mixture  gradually  hardens,  the  calcium  hydrox- 
ide absorbing  carbon  dioxide  and  changing  into  calcium  carbonate. 
The  sand  which  is  added  in  all  probability  serves  the  purpose  of 
rendering  the  mortar  porous,  and  it  thereby  facilitates  the  absorp- 
tion of  carbon  dioxide ;  it  certainly  does  not,  as  was  formerly  sup- 
posed, render  the  product  •  hard  by  forming  calcium  silicate.  A 
mixture  of  quick-lime,  aluminium  oxide,  and  silicon  dioxide  forms 
Portland  cement.  The  latter,  when  brought  in  contact  with  water, 
gradually  hardens,  owing  to  the  union  of  the  calcium  and  alumin- 
ium oxides  with  the  silicon  dioxide  which  is  present.  The  chemi- 
cal process  depends  upon  the  formation  of  hydrated  calcium  silicates 
as  well  as  of  calcium  aluminate. 

Barium  superoxide.  In  addition  to  the  ordinary  oxide,  BaO, 
barium  is  also  able  to  form  a  hyperoxide,  Ba  02  (see  pages  323  and 
324).  This  compound  is  prepared  by  passing  oxygen  over  barium 
oxide  which  is  heated  to  redness,  or  by  heating  a  mixture  of  barium 
oxide  and  potassium  chlorate.  The  substance  is  a  white  powder 
which  loses  oxygen  at  a  bright  red  heat.  It  is  a  powerful  oxidizer : 
hydrogen,  boron,  carbon,  sulphur,  etc.,  are  changed  to  the  corre- 
sponding oxides  when  heated  with  it ;  in  many  cases  the  tempera- 
ture of  the  mass  even  spontaneously  increases  to  redness  during  the 
process.  Barium  hyperoxide,  when  mixed  with  cold  water,  forms  a 
hydrate,  Ba  02  +  6  H2  0  ;  boiling  water  decomposes  it,  liberating 
oxygen  and  leaving  barium  hydroxide.  Barium  hyperoxide  forms 
hydrogen  peroxide  (see  page  50)  when  treated  with  dilute  acids :  — 

Ba02  +  2  HCl=BaCl2  +  H202. 

In  this  respect  barium  hyperoxide  differs  radically  from  other  hy- 
peroxides,  for  the  latter  liberate  chlorine  when  in  contact  with 
hydrochloric  acid  (see  pages  60  and  323).  The  cause  of  this  differ- 
ence lies  in  the  fact  that  the  hyperoxides  of  manganese  and  lead 
form  intermediary  compounds  before  setting  free  the  halogen. 

The  chlorides  of  the  elements  of  this  family  (MC12)  are  all 
soluble  in  water;  those  of  beryllium  and  magnesium  decompose 
when  heated  in  a  current  of  air,  giving  off  chlorine  and  leaving  the 
oxide ;  the  chlorides  of  calcium,  barium,  and  strontium  are  more 
stable.  The  chloride  of  calcium,  CaCl2-|-6  H2  0,  melts  in  its  water 


ALKALINE  EARTHS;   HALIDES,  CARBONATES.  415 

of  crystallization  at  29°  ;  at  100°  it  becomes  anhydrous  and  then 
again  melts;  this  fused  form  of  calcium  chloride  is  deliquescent, 
and,  because  it  greedily  absorbs  moisture,  is  frequently  used  as  a 
drying  agent.  The  chloride  of  strontium  is  not  deliquescent,  while 
the  chloride  of  barium  slowly  takes  up  water  from  the  air.  The 
bromides  and  iodides  are  like  the  chlorides  in  every  respect.  Both 
the  chlorides  of  calcium  and  of  magnesium  occur  as  minerals  in 
some  salt  deposits,  the  former  as  chlorocalcite,  the  latter  as  the 
extremely  deliquescent  mineral  bischofite.  Carnallite  is  a  double 
chloride  of  potassium  and  magnesium,  K  Cl,  Mg  C12  +  6  H2  0,  which 
is  found  quite  frequently  in  the  Strassfurth  salt  regions.  The  in- 
creasing metallic  character  of  the  elements  of  this  family,  as  we 
pass  from  the  member  with  the  smallest  atomic  weight  to  that  with 
the  largest,  is  very  well  illustrated  by  the  behavior  of  the  chlorides 
in  the  presence  of  water :  — 

Be  C12  +  4  H2  O,  completely  decomposed  into  Be  O  +  2  H  Cl  when  its 
solution  is  evaporated. 

Mg  C12  +  6  H2  O,  completely  decomposed  into  Mg  O  +  2  H  Cl  when  its 
solution  is  evaporated. 

Ca  C12  +  6  H2  O,  partly  decomposed  into  basic  calcium  chloride  when  heated 
with  water. 

Sr  C12  +  6  H2  O,  not  decomposed. 

Ba  C12  +  2  H2  O,  not  decomposed. 

All  of  the  carbonates  of  the  members  of  this  family  are  insolu- 
ble in  water.  They  are  the  more  stable  the  more  positive  the  metal 
forming  them  is,  so  that  their  stability  increases  as  the  atomic 
weights  of  the  elements,  counting  from  above  downward,  become 
greater.  Beryllium  carbonate  is  only  capable  of  existence  as  a 
normal  salt  when  it  is  in  an  atmosphere  of  carbon  dioxide ;  when 
exposed  to  the  air  it  breaks  down,  giving  off  carbon  dioxide  and 
leaving  a  basic  carbonate;  magnesium  carbonate  begins  to  break 
down  at  100°,  calcium  carbonate  at  a  low  red  heat,  while  neither 
strontium  nor  barium  carbonate  decomposes  until  a  white  heat  is 
reached.  All  of  the  carbonates  can  be  prepared  by  precipitation 
from  solutions  of  the  salts  of  the  respective  metals  by  addition  of  a 
soluble  carbonate,  such  as  that  of  sodium.  The  following  will  serve 
as  examples :  — 

Na2C03          +BaCl2  =2 Nad     -f  BaCO3 


Soluble.  Insoluble. 


416  ALKALINE    EARTHS;    OCCURRENCE. 

K2C03  +  CaCl2   =  2KC1      +  CaC03 

)2  C03  +  Mg  C12  =  2  NH4  Cl  +  Mg  C03 


Soluble.  Insoluble. 

The  carbonates  of  magnesium,  calcium,  strontium,  and  barium, 
and  the  double  carbonate  of  magnesium  and  calcium,  form  an  ex- 
tremely important  dimorphous  and  isomorphous  group  of  minerals ; 
in  some  localities  entire  mountain  ranges  are  made  up  of  these  com- 
pounds, while  the  various  amorphous  and  cryptocrystalline  varieties 
of  calcium  carbonate  (known  as  chalk,  limestone,  and  marble)  con- 
stitute deposits  of  astonishing  magnitude  (see  page  269).  The  fol- 
lowing table  gives  the  relationship  between  the  crystalline  carbonates 
of  the  elements  belonging  to  this  family :  — 

HEXAGONAL  SYSTEM,  KHOMBOHEDRA.  KHOMBIC  SYSTEM. 

Calcite  group.  Arragonite  group. 

Calcite  (calcspar) Ca  CO3 Arragonite. 

Magnesite Mg  CO3 

Dolomite Ca  CO3  ,  Mg  CO3 

BaCO3 Witherite. 

Sr  CO3 Strontianite. 

While  magnesium  carbonate  is  not,  by  itself,  capable  of  crystallizing 
in  the  same  form  as  arragonite,  yet,  when  mixed  with  the  carbonate 
of  manganese  or  of  calcium,  it  assumes  that  form ;  on  the  other 
hand,  calcite  crystals  which  contain  barium  and  strontium  are  also 
found.  This  isodimorphous  group  of  minerals  is  very  far  reaching, 
for"  the  carbonates  of  zinc,  iron,  manganese  and  cobalt  also  belong  to 
the  calcite  group,  while  those  of  manganese,  iron  and  lead  are  also 
found  crystallizing  in  the  form  of  arragonite. 

The  sulphates  of  the  elements  of  this  family  may  be  divided  into 
two  classed,  those  of  beryllium  and  magnesium,  ^which  are  soluble 
in  water,  and  those  of  calcium,  strontium,  and  barium,  which  are  in- 
soluble, or  nearly  insoluble.  Of  the  latter  class,  that  of  calcium  is 
soluble  with  difficulty,  that  of  strontium  is  less  soluble  than  that  of 
calcium,  while  the  sulphate  of  barium  is  insoluble ;  these  sulphates 
can  therefore  be  produced  as  white  precipitates  on  the  addition  of 
a  soluble  sulphate  to  the  solutions  of  salts  of  the  respective  metals.* 

*  Calcium  sulphate  will  not  be  precipitated  if  the  solution  of  the  calcium 
salt  is  too  dilute.  A  solution  of  calcium  sulphate  will  precipitate  strontium 
and  barium  salts;  a  solution  of  strontium  sulphate  will  precipitate  barium 
salts. 


THE   VITKIOLS. 


417 


Magnesium  sulphate  is  the  representative  of  a  large  number  of 
sulphates  which  are  known  as  vitriols.  The  vitriols,  with  the 
exception  of  copper  sulphate  (which  contains  five  molecules  of 
water)  all  crystallize  with  seven  molecules  of  water  of  crystalliza- 
tion and  form  a  typical  isomorphous  group  of  compounds  ;  the  one 
exception,  copper  sulphate,  Cu  S04  +  5  H2  0,  can,  however,  crystal- 
lize with  seven  molecules  of  water  when  it  is  present  in  an  isomor- 
phous mixture  with  some  other  vitriol.  The  vitriols,  when  heated  to 
100°,  change  to  salts  having  the  composition  MS04  +  H2  0 ;  the 
last  remaining  molecule  of  water  passes  off  at  a  higher  temperature, 
and  for  this  reason  these  substances  are  commonly  regarded  as 
being  secondary  salts  of  the  hydrated  sulphuric  acid,  H4  S05 ,  so 
that  their  formulae  would  be  MH2  S05  +  6  H2  0.  The  following  is 
a  list  of  the  compounds  comprising  this  isomorphous  group :  — 

Be  H2  SO5  +  6  H2  O Beryllium  sulphate. 

Mg  H2  SO5  +  6  H2  O Magnesium  sulphate. 

ZnH2SO5 +6H2O. Zinc  sulphate  (white  vitriol). 

FeH2SO5+6H2O Ferrous  sulphate  (green  vitriol). 

Ni  H2  So5  +6H2O Nickel  sulphate. 

Co  H2SO5  +  6H2O. Cobalt  sulphate. 

CuH2SO6  +  4H2O Copper  sulphate  (blue  vitriol). 

Copper  sulphate  crystallizes  with  six  molecules  of  water  when  in  an  iso- 
morphous mixture  with  one  or  more  of  the  other  vitriols.  All  of  the  vitriols 
can  crystallize  with  one  formula  weight  of  potassium  or  of  ammonium  sulphate 
to  form  double  salts  of  the  general  formula  M'2  SO4 ,  M"SO4  +  6  H2  O. 

The  sulphates  of  calcium,  strontium,  and  barium  occur  quite  fre- 
quently as  minerals ;  they  constitute  an  isomorphous  group,  which 
crystallizes  in  the  rhombic  system  ;  the  representatives  of  this  group 
are  given  in  the  following  table  :  — 


NAME  OF  MINERAL. 

FORMULA. 

1.  Anhydrite 

1. 

CaSO4 

2.  Celestine 

2. 

SrS04 

3.  Barite 

3. 

BaSO4 

4.  Barytocelestine 

4. 

(Ba,  Sr)SO4 

5.  Anglesite 

5. 

PbSO4 

Calcium  sulphate  also  occurs  as  gypsum,  in  which  mineral  it 
crystallizes  with  two  molecules  of  water  of  crystallization,  Ca  S04 
-j-  2  H2  0 ;  gypsum  is  frequently  found  as  a  massive  variety,  in 


418  CALCIUM    SULPHATE,    PHOSPHATES. 

which  condition  it  bears  the  name  of  alabaster.  When  gypsum  is 
heated  to  a  little  above  100°  it  loses  its  water  of  crystallization, 
and  is  converted  into  a  fine  white  powder  which  is  termed  plaster 
of  Paris ;  this  substance,  when  mixed  with  water,  once  more  unites 
with  that  liquid,  and  then  changes  into  a  firm  mass  which  has  the 
composition  of  alabaster;  this  process  is  termed  the  "  setting"  of 
plaster  of  Paris.  Many  natural  waters  contain  calcium  sulphate  in 
solution  ;  such  waters  are  known  as  permanent  hard  waters,  because 
the  calcium  salt  is  not  removed  from  them  by  boiling.  Temporary 
hard  waters  contain  primary  calcium  carbonate,  Ca  (HC03)2;  the 
latter  substance,  however,  breaks  down  into  carbon  dioxide  and 
secondary  calcium  carbonate  (insoluble)  when  the  solution  is  boiled  ; 
the  calcium  which  is  contained  in  the  primary  carbonate  is,  there- 
fore, entirely  removed  by  this  process.  Calcium  is  able  to  form  an 
insoluble  compound  with  soap ;  hard  waters,  therefore,  form  a  pre- 
cipitate when  brought  in  contact  with  the  latter  substance,  and  will 
consequently  form  a  lather  only  after  all  the  calcium  salts  have 
been  removed  (see  page  43). 

The  tertiary  and  secondary  phosphates  of  the  elements  belonging 
to  this  family  are  all  insoluble  in  water ;  the  primary  phosphates 
are  soluble ;  as  a  consequence,  the  tertiary  phosphates  are  dissolved 
on  addition  of  the  stronger  acids,  such  as  hydrochloric  or  nitric. 
The  tertiary  phosphate  of  calcium  is  the  only  one  of  these  which 
occurs  as  a  mineral,  the  name  of  which  is  osteolite,  Ca3(P04)2; 
this  substance  is  often  found  in  massive  deposits,  which  are  espe- 
cially extensive  in  Florida ;  guano  is  tertiary  calcium  phosphate 
which  is  mixed  with  a  number  of  impurities,  such  as  calcium  car- 
bonate, magnesium  carbonate,  gypsum,  etc.  A  double  salt  of  cal- 
cium phosphate  and  calcium  chloride  is  found  in  crystals  belonging 
to  the  hexagonal  system,  and  is  termed  apatite,  Ca3  (P04)2  -f  Ca  C12 . 
In  addition  to  its  occurrence  as  mineral  deposits,  calcium  phosphate 
is  the  chief  constituent  of  the  inorganic  portions  of  the  bones  (see 
page  211).*  The  reactions  relating  to  the  conversion  of  the  terti- 
ary phosphate  into  the  soluble  primary  one  are  given  on  pages  229 
and  230.  The  soluble  primary  phosphate  of  calcium,  mixed  with 
gypsum  and  other  impurities,  goes  by  the  name  of  superphosphate ; 
this  substance  is  used  as  a  foundation  for  the  mixtures  which  find 
their  way  into  the  market  as  artificial  fertilizers.  Of  course,  the 

*  Bone-ash  contains  85  per  cent  of  calcium  phosphate. 


MAGNESIUM   PHOSPHATE.  419 

phosphate  in  a  fertilizer  must,  in  part  at  least,  be  in  a  soluble 
form  so  that  it  can  be  readily  absorbed  by  plants  ;  the  conversion 
of  the  insoluble  tertiary  into  the  soluble  primary  phosphate  is 
effected  by  means  of  sulphuric  acid  :  — 


The  value  of  a  fertilizer  is  determined  by  the  amount  of  soluble 
phosphate  which  it  contains. 

Magnesium,  phosphate  behaves  exactly  as  does  calcium  phos- 
phate, the  tertiary,  secondary,  and  primary  phosphate  being  known  ; 
the  latter  is  soluble  in  water.  The  most  important  phosphate  of 
magnesium  is  the  insoluble  ammonium-magnesium  phosphate, 
Mg(NH4)P04;  the  formation  of  this  salt  as  a  precipitate  may  be 
used  as  a  test  for  the  presence  either  of  magnesium  or  of  phos- 
phoric acid  in  a  solution.  The  salt  is  produced  by  adding  a  soluble 
phosphate  to  a  mixture  of  a  soluble  magnesium  salt  with  ammonia 
and  ammonium  chloride.  If  ammonia  alone  were  added  to  a  solu- 
tion containing  a  magnesium  salt,  a  portion  of  the  latter  would  be 
decomposed  and  the  base  precipitated  as  magnesium  hydroxide, 
while  a  part  would  remain  in  solution  as  a  double  salt  of  magnesium 
and  ammonium,  for  magnesium  salts  have  the  power  of  forming 
compounds  with  ammonium  salts,  and  these  compounds  are  not 
decomposed  by  ammonia  :  — 

2  Mg  S04  +  2  NH3  +  2  H2  0  =  Mg  (OH  )2  +  (  KH4)2  S04  ,  Mg  S04  . 

The  previous  addition  of  an  ammonium  salt  to  a  solution  containing 
a  salt  of  magnesium  therefore  prevents  any  precipitation  of  mag- 
nesium hydroxide  by  means  of  ammonia.  When  ammonium-mag- 
nesium phosphate  is  heated,  it  loses  ammonia  and  changes  into  the 
secondary  phosphate  of  magnesium  :  — 

MgKH4  P04  =  Mg  HP04  +  NH,  , 

and  the  latter,  on  further  heating,  again  loses  water,  and  finally 
leaves  magnesium  pyrophosphate  (see  page  231)  :  — 

2  Mg  HP04  =  Mg2P2  07  +  H2  0. 

Arsenic  acid  reacts  similarly  to  phosphoric  acid.  It  forms  an 
ammonium-magnesium  arsenate,  MgKB4As04.  The  precipitate 
formed  is  not  to  be  distinguished  from  ammonium-magnesium  phos- 
phate ;  if  both  arsenic  and  phosphoric  acid  are  present  in  any  solu- 


420  CALCIUM  SILICATES. 

tion,  then  the  arsenic  acid  must,  previous  to  the  precipitation  of 
the  phosphoric  acid,  be  reduced  to  arsenious  acid  by  means  of  sul- 
phur dioxide,  for  in  this  form  it  is  not  precipitated  by  the  mixture 
of  magnesium  salts. 

The  hypochlorite  of  calcium,  when  mixed  with  calcium  hydroxide 
and  calcium  chloride,  is  known  as  chloride  of  lime.  Because  of 
the  invariable  occurrence  of  calcium  chloride  in  conjunction  with 
calcium  hypochlorite,  the  theory  is  not  unfrequently  maintained 
that  calcium  hypochlorite  is  in  reality  a  mixed  chloride  and  hypo- 
chlorite of  calcium  having  the  formula  :  — 


Ca 


(0( 
JC1. 


C1 


Such  a  body  would  have  the  same  composition  by  weight  as  an 
equimolecular  mixture  of  calcium  chloride  and  calcium  hypochlo- 
rite, Ca  C12  +  Ca  (0  Cl)2  =  2  Ca  (0  Cl)  Cl.  The  reactions  peculiar 
to  calcium  hypochlorite  have  been  explained  on  pages  122  and  123. 

The  chlorates  and  nitrates  of  calcium,  barium,  and  strontium  are 
soluble  in  water  and  are  extensively  used  in  the  manufacture  of 
Greek  fire  ;  the  chlorate  of  strontium,  when  mixed  with  oxidizable 
substances  and  ignited,  gives  an  intensely  red  light,  while  that  of 
barium  produces  a  green  one. 

The  silicates  of  calcium  are  of  the  greatest  importance  because 
of  the  fact  that  they  are  essential  in  the  manufacture  of  glass. 
Calcium  metasilicate,  CaSi03,  occurs  as  the  mineral  wollastonite, 
and  a  great  many  other  naturally  occurring  silicates  contain  cal- 
cium (see  page  307)  ;  these  silicates  are,  however,  crystalline  in 
their  structure,  while  the  artificial  silicates  are,  as  a  rule,  amor- 
phous. Glass  consists  of  a  vitreous  mixture  of  the  silicates  of  the 
alkalies  and  of  calcium,  with  silicon  dioxide;  in  some  forms  of 
glass,  however,  lead  oxide  may  replace  calcium  oxide.  Ordinary 
window  glass  is  produced  by  fusing  sand,  calcium  carbonate,  and 
sodium  carbonate  together;  the  silicate  of  calcium  and  sodium  so 
formed  is,  in  reality,  of  a  crystalline  structure,  but  this  structure 
is  concealed  by  the  vitreous  mass  of  silicon  dioxide ;  such  a  glass 
is  readily  attacked  by  laboratory  reagents  or  even  by  the  continued 
action  of  water,  and,  after  it  has  been  acted  on  for  some  time,  the 
crystalline  condition  becomes  apparent.*  Window  glass  is  first 

*  All  kinds  of  glass  are  more  or  less  attacked  by  water,  or  alkalies;  dilute 
acids  do  not  appear  to  have  any  effect ;  concentrated  acids  prevent  the  solution 


GLASS.  421 

blown  and  then  cut  into  suitable  pieces  ;  for  that  reason  it  is  more 
or  less  irregular  in  thickness  and  does  not  present  a  perfectly 
smooth  surface.  Plate  glass  has  essentially  the  same  composition, 
but  is  cast  on  flat  plates  and  finally  polished.  Bohemian  glass  is 
made  by  fusion  of  potassium  carbonate,  a  little  sodium  carbonate, 
silicon  dioxide,  and  calcium  carbonate.  The  replacing  of  the 
sodium  carbonate  by  potassium  carbonate,  with  the  resulting  for- 
mation of  a  potassium-calcium  silicate,  renders  the  glass  difficult 
to  fuse.  Bohemian  glass  is  used  for  the  manufacture  of  chemical 
apparatus,  a  further  advantage  belonging  to  this  variety  being  in 
the  fact  that  it  is  not  readily  attacked  by  chemical  reagents. 
Bottle  glass  contains  more  calcium  silicate  than  either  of  the 
above  varieties  ;  it  is  frequently  colored  green  by  the  presence 
of  ferrous  silicate.  The  commoner  bottles  are  made  from  impure 
materials.  Flint  glass  is  of  similar  composition  to  ordinary  lime- 
soda  glass,  with  the  exception  that  the  lime  is  replaced  by  lead 
oxide  ;  it  is  characterized  by  having  a  very  high  index  of  refrac- 
tion,* great  lustre  and  high  specific  gravity  ;  it  is  the  most  fusible 
variety  of  glass.  |  Flint  glass  is  used  in  the  manufacture  of  opti- 
cal instruments  and  in  some  chemical  apparatus.  Strass  is  flint 
glass  which  is  very  rich  in  lead;  it  is  used  for  making  artificial 
gems.  Glass  is  stained  by  adding  inorganic  coloring  matter  to  the 

of  the  glass  to  a  certain  extent.  The  solubility  of  glass  in  solutions  of  sodium 
hydroxide  or  sodium  carbonate  (containing  -fa  of  a  grain-molecular  weight 
in  one  litre  of  water)  is  about  three  to  five  times  the  solubility  of  the  same 
glass  in  pure  water.  This  relationship  varies  with  different  kinds  of  glass. 
The  solubility  is  greater  in  sodium  carbonate  than  it  is  in  sodium  hydroxide 
solution.  A  glass  which  is  made  up  as  follows  has  the  best  composition  for 
general  laboratory  purposes  :  — 

K2    0 6.2  per  cent.  Mn  O 0.2  per  cent. 

Na2  O 6.4  per  cent.  A1.2  O3    .."...     0.4  per  cent. 

Ca  O 10.0  per  cent.  Si    O2 76.8  per  cent. 

Such  a  glass  loses  .0037  gram  per  100  sq.  c.  m.  to  sodium  hydroxide  solution 
of  the  above  strength  and  .0059  gram  per  sq.  c.  m.  to  sodium  carbonate.  (See 
Foerster;  Berichte  d.  Deutsch.  Chem.  Gesell.;  26,  2915.) 

*  The  index  of  refraction  for  flint  glass  is  1.8,  while  that  of  window  glass 
is  1.53. 

t  In  working  with  lead  glass  care  must  be  taken  not  to  bring  the  same 
into  the  reducing  flame  (see  page  283),  which  is  that  portion  immediately  out- 
side of  the  central  zone,  otherwise  a  part  of  the  lead  silicate  will  be  reduced 
and  lead  will  separate ;  the  latter  renders  the  glass  black  and  opaque. 


422  GLASS. 

colorless  varieties  and  fusing;  thus,  blue  glass  is  produced  by  add- 
ing a  little  cobalt  salt,  green  glass  by  copper  and  chromium,  etc. 
The  various  glass  utensils  which  are  used  must  be  previously  an- 
nealed by  a  very  slow  cooling  process ;  if  this  precaution  is  not 
taken,  the  outer  surfaces,  cooling  more  rapidly  than  the  remainder 
of  the  mass,  establish  such  a  tension  that  the  slightest  scratch  upon 
the  surface  will  cause  the  entire  object  to  be  shattered.  This  con- 
dition is  best  shown  in  the  so-called  Prince  Rupert's  drops.  The 
latter  are  made  by  fusing  glass  and  allowing  the  drops  to  fall  into 
water;  when  the  end  of  the  small  pear-shaped  mass  so  formed  is 
broken,  the  drop  disintegrates  into  a  sandy  mass  with  explosive 
violence. 


ZLNTC;  CADMIUM;  MEECUKY. 


423 


CHAPTER   LV. 

ZINC,   CADMIUM,  AND    MERCURY. 

Zinc  ;  symbol,  Zn ;   atomic  weight,  65.3  ; 
Cadmium  ;  symbol,  Cd ;   atomic  weight,  112  ; 
Mercury  ;  symbol,  Hg ;    atomic  weight,  200. 

ZINC,  cadmium,  and  mercury  form  the  secondary  group  of  the 
family  of  the  alkaline  earths.  They  are  the  second  elements  of 
the  second  halves  of  the  long  periods,  while  calcium,  strontium,  and 
barium  occupy  the  same  position  in  the  first  halves.  As  a  conse- 
quence they  bear  much  the  same  relationship  to  the  alkaline  earths 
as  copper,  silver,  and  gold  bear  to  the  alkalies.  The  typical  oxides 
and  hydroxides,  as  well  as  the  salts,  derived  from  zinc,  cadmium,  and 
mercury,  are,  therefore,  of  the  same  formula  as  they  are  with  cal- 
cium, strontium,  and  barium.  In  both  divisions  of  the  family  the 
metals  are  divalent,  so  that  the  oxides  have  the  general  formula 
MO,  and  the  hydroxides  M  (OH)2 

In  their  physical  characteristics  the  three  elements  are  all,  most 
certainly,  metallic  in  their  nature ;  but  zinc,  the  one  with  the  least 
atomic  weight,*  is  less  positive  than  the  other  two.  The  change  in 
physical  character  brought  about  by  the  increasing  atomic  weight, 
as  we  pass  from  zinc  to  mercury,  is  shown  in  the  following  table :  — 


METAL. 

ATOMIC  WEIGHT. 

SPECIFIC  GRAVITY. 

ATOMIC  VOLUME. 

MELTING  POINT. 

Zinc 

65.3 

7.15 

9.1 

417° 

Cadmium 

112. 

8.65 

12.9 

317° 

Mercury 

200. 

13.59 

14.7 

—  39° 

All  three  of  the  metals  are  volatile ;  their  boiling  points  decrease 
with  increasing  atomic  weight,  just  as  their  melting  points  do ;  this 

*  Zinc  is  also  the  element  with  the  next  lower  atomic  weight  to  that  of 
gallium,  which  element  has  many  of  the  characteristics  of  a  not-metal;  we 
should  therefore  scarcely  expect  zinc  to  present  very  marked  metallic  properties. 


424  ZINC  ;    PROPERTIES. 

phenomenon,  which  is  most  strikingly  illustrated  in  the  case  of  the 
trio  of  metals  under  discussion,  is  exactly  the  reverse  of  the  changes 
in  the  melting  points  taking  place  in  the  not-metallic  families  at  the 
right  hand  extremities  of  the  periods.  It  will  be  remembered,  also, 
that  the  melting  points  of  the  alkali  metals  diminish  as  we  pass 
from  member  to  member  in  the  direction  of  increasing  atomic 
weights,  and  the  same  is  true  of  the  metals  constituting  the  first 
portion  *  of  the  family  under  discussion ;  the  elements,  therefore, 
which  comprise  the  first  two  families  in  the  periodic  system  show 
decreasing  melting  points  with  increasing  atomic  weights ;  whether 
the  same  is  true  of  the  boiling  points  cannot  be  stated,  as  many  of 
the  elements  cannot  be  volatilized  by  any  of  the  means  at  our  com- 
mand ;  it  certainly  is  true  of  the  three  metals  under  consideration, 
for:  — 

Zinc  boils  at  927°. 

Cadmium  boils  at  772°. 

Mercury     boils  at  357°. 

Zinc  is  an  element  with  a  brilliant  metallic  lustre  which  pos- 
sesses a  bluish  tint  and  a  crystalline  structure,  f  It  is  malleable 
only  at  temperatures  between  100°  and  150° ;  at  ordinary/ 
temperatures  it  is  easily  fractured ;  at  200°  the  metal  can  readily 
be  pounded  into  a  powder;  when  heated  to  its  point  of  vaporization 
in  the  air,  zinc  burns  with  a  bluish  white  flame,  producing  zinc 
oxide,  Zn  O.  The  metal  easily  dissolves  in  dilute  acids  (see  pages 
33,  206).  Solutions  of  the  caustic  alkalies,  when  warmed,  attack 
zinc,  forming  zincates  and  liberating  hydrogen ;  in  this  respect  zinc 
resembles  aluminium  and  a  number  of  other  metals  which  can  dis- 
play both  metallic  and  not-metallic  properties ;  this  behavior  is  not 
unexpected  when  we  consider  that  zinc  is  the  element  with  least 
metallic  properties  in  the  group  we  are  considering,  and  that  every 
other  element  following  it  in  the  same  period  can  also  display  certain 
not-metallic  properties. 

Cadmium  is  a  glistening,  tin-like  metal ;  it  is  .  soft,  though 
harder  than  tin,  and  it  has  a  crystalline  structure ;  when  heated  to 

*  Beryllium,  magnesium,  calcium,  strontium,  and  barium  have  decreasing 
melting  points  as  we  pass  along  the  line  in  the  order  named.  The  melting 
points  of  calcium  and  strontium  are  not  accurately  determined. 

t  When  pieces  of  zinc  are  bent,  a  peculiar  crepitation,  similar  to  that  with 
crystallized  tin,  is  observed  (see  page  313). 


CADMIUM  ;    MERCURY ;    PROPERTIES.  425 

its  boiling  point  in  the  air,  it  burns  like  zinc,  forming  cadmium 
oxide,  Cd  0.  The  metal  dissolves  in  acids  less  readily  than  does 
zinc,  but  with  this  exception  shows  the  same  behavior.  The  metal 
does  not  dissolve  in  solutions  of  the  caustic  alkalies. 

Mercury  is  a  bluish  white  metal  which  is  fluid  under  ordinary 
circumstances.  The  solid  substance  (formed  at  —  39°)  is  soft  and 
malleable  when  pure.  Mercury  is  not  changed  in  the  air  at  ordi- 
nary temperatures ;  if  heated  for  some  time  near  its  boiling  point 
and  in  the  presence  of  oxygen,  it  is  changed  to  red  mercuric  oxide 
(  Hg  0)  ;  ozone  readily  attacks  it  without  the  necessity  of  heating. 
Hydrochloric  acid,  or  dilute  sulphuric  acid,  does  not  dissolve  mer- 
cury ;  concentrated  sulphuric  acid,  when  hot,  attacks  the  metal  and 
liberates  sulphur  dioxide  ;  nitric  acid,  concentrated  or  dilute,  acts 
upon  mercury ;  if  the  acid  is  dilute,  mercurous  nitrate  is  formed,  if 
concentrated,  mercuric  nitrate  is  produced ;  alkalies  do  not  attack 
mercury. 

As  has  been  stated,  zinc,  cadmium,  or  mercury  can  be  easily  vol- 
atilized. The  specific  gravities  of  their  vapors  are  as  follows  :  — 

Zinc,  specific  gravity  of  vapor,  air  =  1,  2.36,  H2  =  2,   67.96;  molecular  weight,    65.3. 

Cadmium,  specific  gravity  of  vapor,  air  =  1,  3.95,  H2  =  2, 113.47;  molecular  weight,  112. 
Mercury,     specific  gravity  of  vapor,  air  =  1,  6.83,  H2  =  2, 196.7;     molecular  weight,  200. 

The  above  determinations  show  that  the  molecular  weight  and 
the  atomic  weight  of  each  of  the  three  elements,  when  they  are  in 
the  state  of  a  vapor,  are  identical.  Zinc,  cadmium,  and  mercury, 
therefore,  in  changing  from  the  liquid  to  the  gaseous  state,  separate 
at  once  into  the  individual  atoms,  provided  our  decision,  as  to  what 
the  relative  weights  of  these  atoms  are,  is  the  correct  one ;  the 
only  other  elements  which  exhibit  the  same  phenomenon  are  so- 
dium and  potassium  (and  iodine  when  heated  above  1600° ;  on  de- 
creasing that  temperature  the  atoms  of  iodine  gradually  unite  to 
form  diatomic  molecules  [see  page  84]  )  .  In  the  course  of  our  study 
we  have  therefore  encountered  elements  with  one,  with  two,  with 
three,  and  with  four  atoms  united  to  form  a  molecule  of  the  gas ;  the 
molecules  having  three  and  four  atoms  are,  however,  dissociated  at 
high  temperatures  and  then  change  into  those  having  two,  while  some 
of  the  latter  have  already  been  dissociated  into  the  individual  atoms. 
Undoubtedly,  were  we  able  to  command  a  sufficiently  high  tempera- 
ture in  the  apparatus  used  for  determining  the  vapor  densities,  we 
should  be  able  to  discover  that  all  diatomic  molecules  can  be  changed 


426 


ZINC  :    CADMIUM  :    OCCURRENCE. 


into  monatomic  ones.  The  following  gives  a  list  of  elements  of 
which  it  has  been  possible  to  obtain  specific  gravities  while  they 
were  in  the  gaseous  state ;  in  the  cases  of  selenium  and  antimony 
some  doubt  exists  as  to  whether  the  molecules  are  in  reality  tria- 
tomic,  or  whether  the  vapor  density  numbers  which  have  been 
obtained  are  only  accidental :  — 


MONATOMIC.* 

DIATOMIC. 

TKIATOMIC. 

TETKATOMIC. 

Sodium 

Hydrogen 

Ozone 

Phosphorus 

Potassium 

Chlorine 

Selenium  (?) 

Arsenic 

Zinc 

Bromine 

Antimony  (?) 

Cadmium 

Iodine  (below  600°) 

Mercury 

Oxygen 

Iodine  (above  1600°) 

Sulphur  (above  1000°) 

Bismuth  (?) 

Selenium 

Tellurium  (?) 

Nitrogen 

Phosphorus    f    ,a? 
4  white 
Arsenic          j  heat 

The  fact  that  so  many  elementary  gases  are  formed  of  complex 
molecules  was  not  understood  when  the  theory  that  equal  volumes 
of  gases  contain  equal  numbers  of  particles  was  first  advanced ;  the 
discrepancy  frequently  observed  between  the  specific  gravity  of 
the  gases  and  the  atomic  weights  of  elements  determined  by  other 
means,  therefore,  led  to  a  disbelief  in  Avogadro's  hypothesis,  and  to 
considerable  confusion  in  the  determination  of  atomic  weights. 

The  principal  minerals  in  which  zinc,  cadmium,  and  mercury 
occur  are  as  follows  :  — 

Zinc  and  cadmium.  The  occurrence  of  uncombined  zinc  as  a 
mineral  is  doubtful.  Zinc  and  cadmium  occur  as  the  sulphides, 
Zn  S  and  Cd  S,  in  an  isomorphous  and  dimorphous  group  which  also 
includes  the  sulphides  of  manganese,  iron,  and  nickel.  The  sulphide 
of  zinc  occurs  in  crystals  belonging  to  the  regular  system  termed 
sphalerite  or  zinc-blende,  and  in  hexagonal  crystals  (wurtzite).  The 
sulphide  of  cadmium  is  isomorphous  with  wurtzite,  and  is  known  as 
greenockite.  Zinc  is  sometimes  found  as  zincite,  which  is  the  ox- 

*  It  seems  scarcely  necessary  to  state  that  the  term  monatomic  means  ex- 
isting as  molecules  formed  of  one  atom,  diatomic  of  two,  etc. 


ZINC;  CADMIUM;  MERCURY;  METALLURGY.  427 

ide,  ZnO,  colored  red  by  means  of  manganese.  The  carbonate  of 
zinc,  Zn  C03  (smithsonite),  is  isomorphous  with  calcite  (see  page 
416).  An  aluminate  of  zinc,  Zn  (Al  02),  isomorphous  with  spinell 
(see  page  339),  is  also  met  with. 

Mercury  is  sometimes  found  in  small,  fluid  globules  in  places 
where  the  most  important  mineral  containing  mercury  (namely,  the 
red  sulphide,  cinnabar  [HgS],)  also  occurs.  In  addition,  amal- 
gams of  mercury  with  silver  and  gold  are  sometimes  met  with. 

Zinc  is  obtained  from  its  ores  by  roasting  the  sulphide  in  a 
draught  of  air,  and  by  subsequently  heating  with  charcoal  the  oxide 
produced  by  this  means.  As  zinc  is  volatile  at  no  very  high  tem- 
perature, the  mixture  of  oxide  and  charcoal  is  placed  in  earthen- 
ware retorts  which  are  gradually  heated ;  carbon  monoxide  passes 
off  during  this  process  :  — 

Zn  0  -f  C  =  Zn  +  CO. 

Finally,  the  temperature  is  increased  to  a  point  at  which  zinc  begins 
to  distil;  earthenware  receivers  are  then  placed  before  the  open- 
ings of  the  retorts,  and  the  metal  is  collected  therein.  The  first 
portions  which  pass  over  are  deposited  on  the  walls  of  the  receivers 
in  the  form  of  a  fine  dust  which  always  contains  more  or  less  oxide 
of  zinc ;  this  product,  which  is  known  as  zinc-dust,  is  very  fre- 
quently used  as  a  reducing  agent  in  the  laboratory.  The  fused  zinc 
which  finally  collects  is  generally  impure,  containing  lead,  iron,  and 
cadmium ;  it  is  separated  from  those  metals  by  repeated  distillation. 
The  cadmium,  having  a  lower  boiling  point,  passes  over  first,  while 
the  lead  and  iron  remain  behind.  The  preparation  of  cadmium  is 
like  that  of  zinc.  As  cadmium  ores  generally  contain  zinc,  the 
metal  is  separated  from  the  latter  by  distillation.  Mercury  is  ob- 
tained by  roasting  the  sulphide,  the  mercury .  which  passes  off  being 
collected  in  receivers  which  are  connected  with  the  oven.  The  ad- 
dition of  charcoal  is  unnecessary  during  this  process,  because  the 
oxide  of  mercury,  which  would  be  formed  by  roasting  the  sulphide, 
is  further  decomposed  into  mercury  and  oxygen  by  heat  (see  page 
18). 

As  has  been  mentioned,  zinc  is  a  constituent  of  the  alloy  known 
as  brass.  When  sheet-iron  is  coated  with  zinc,  it  is  known  as  gal- 
vanized iron.  Zinc  readily  forms  an  amalgam  with  mercury ;  an 
extensive  commercial  use  of  this  fact  is  made  in  amalgamating  zinc 


428  ZINC;  CADMIUM;  OXIDES. 

battery  plates,  the  latter  being  cleaned,  dipped  in  acid,  and  rubbed 
with  mercury  so  as  to  produce  a  thin  layer  of  amalgam.  Cadmium, 
zinc,  and  mercury  form  an  amalgam  which  readily  hardens ;  this  sub- 
stance is  used  in  filling  teeth.  Alloys  of  cadmium  with  lead  and  bis- 
muth are  used  where  a  very  low-melting  metal  is  required.  Alloys 
of  mercury  are  termed  amalgams ;  a  number  of  these  have  definite 
composition  and  crystalline  form.  The  nature  of  amalgams  has 
been  discussed  on  page  250. 

Zinc  and  cadmium  form  but  one  oxide  apiece ;  these  oxides,  in 
formulae,  correspond  to  the  typical  oxide  of  the  family,  M  0.  The 
hydroxides,  M(OH)2,  can  be  produced  by  adding  soluble  hydrox- 
ides to  solutions  containing  salts  of  cadmium  or  zinc :  — 

MC12  +  2KOH  =  2KC1  +  M(OH)2. 

However,  in  adding  a  caustic  alkali  to  a  zinc  salt,  care  must  be 
taken  not  to  use  an  excess  of  the  reagent ;  for  zinc  hydroxide  acts 
as  an  acid  when  in  the  presence  of  strong  bases,  dissolving  in  the 
latter  to  form  zincates :  — 

Zn(OH)2  +  2KOH  =  Zn  (OK)2  +  2 H2 0, 

(see  pages  315  and  322)  ;  the  zincate  so  formed  is,  however,  de- 
composed by  boiling,  zinc  hydroxide  being  precipitated ;  ammonia 
water  has  the  same  effect  as  solutions  of  the  caustic  alkalies,  an 
excess  of  that  reagent  dissolving  the  precipitated  hydroxide  while 
producing  ammonium  zincate.*  Acid  solutions  of  zinc  salts,  or 
mixtures  of  ammonium  and  zinc  salts,  are  not  affected  by  ammonia. 
In  the  case  of  cadmium,  caustic  alkalies  precipitate  the  hydroxide, 
which,  however,  is  not  dissolved  by  an  excess  of  the  reagent,  cad- 
mium hydroxide  having  no  acid  properties  ;  on  the  other  hand, 
ammonia,  when  in  the  presence  of  ammonium  salts,  produces  no 
precipitate,  for  cadmium  resembles  zinc  and  magnesium  in  the  facil- 
ity with  which  its  salts  form  compounds  with  those  of  ammonium. 

Zinc  oxide  is  a  white  powder ;  yellow  when  heated. 

Cadmium  oxide  is  a  brown  powder. 

Zinc  hydroxide  is  a  white  powder;  changes  to  zinc  oxide  and 
water  when  heated. 

*  Difference  between  zinc  and  magnesium,  for  ammonia  precipitates  a 
portion  of  the  magnesium  as  magnesium  hydroxide.  The  latter  is  insoluble 
in  an  excess  of  the  reagent  (see  page  419). 


ZINC  ;   CADMIUM  ;    SALTS   OF.  429 

Cadmium  hydroxide  is  a  white  powder ;  changes  to  cadmium 
oxide  and  water  when  heated. 

In  addition  to  the  methods  for  the  preparation  of  the  oxides 
which  have  been  given  above,  those  compounds  can  also  be  pro- 
duced by  heating  the  respective  carbonates  or  nitrates. 

The  chloride  of  zinc  can  be  produced  by  heating  zinc  in  a  cur- 
rent of  chlorine,  or  by  dissolving  zinc,  or  the  oxide  or  hydroxide  of 
the  metal,  in  hydrochloric  acid,  evaporating  the  solution  and  dis- 
tilling. The  fused  salt  is  cast  into  sticks  which  are  extremely 
deliquescent.  The  salt  crystallizes  from  a  concentrated  solution  of 
hydrochloric  acid  in  crystals  having  the  formula  Zn  C12  +  H2  0 ; 
but  when  an  aqueous  solution  of  zinc  chloride  is  heated,  a  partial 
decomposition  into  the  basic  chloride  takes  place :  — 

Zn 

and  this  substance,  when  boiled  with  water,  finally  loses  all  chlo- 
rine and  changes  to  the  hydroxide  (see  magnesium  chloride,  pages 
414,  415).  The  chloride  of  cadmium,  Cd  C12 ,  is  similar  to  that  of 
zinc  ;  like  zinc  chloride,  it  is  volatile,  but  is  not  decomposed  when 
the  solution  is  evaporated. 

Zinc  sulphate,  formed  by  dissolving  either  the  hydroxide,  oxide, 
carbonate,  or  the  metal  in  sulphuric  acid,  belongs  to  the  class  of  sul- 
phates which  are  termed  vitriols  (see  page  417).  Like  all  of  the 
sulphates  belonging  to  this  group,  it  is  soluble  in  water  and  crystal- 
lizes with  seven  molecules  of  water,  six  of  which  it  loses  at  100°, 
while  the  seventh  passes  off  only  at  a  higher  temperature. 

Cadmium,  sulphate  does  not  belong  to  che  group  of  vitriols.  Its 
crystals  have  the  formula  3  Cd  S04  -j-  8  H2  0. 

When  a  solution  of  the  carbonate  of  .an  alkali  metal  or  of 
ammonium  is  added  to  the  solution  of  a  zinc  salt,  an  insoluble  basic 
carbonate  of  zinc  is  precipitated.  This  substance  has  a  varying 
composition,  according  to  the  conditions  under  which  it  is  produced. 
The  normal  secondary  carbonate  of  zinc,  Zn  C03 ,  occurs  as  the  min- 
eral smithsonite,  belonging  to  the  calcite  group.  The  carbonate 
is  easily  decomposed  into  zinc  oxide  and  carbon  dioxide  when  it  is 
heated. 

Cadmium  carbonate  is  precipitated  from  solutions  containing  a 
cadmium  salt  by  addition  of  a  soluble  carbonate.  Owing  to  the 


430  MERCURY;  OXIDES  OF. 

more  metallic  nature  of  cadmium,  the  precipitate  so  formed  con- 
sists of  the  secondary  carbonate,  Cd  C03  . 

Zinc  sulphide  is  precipitated  from  the  neutral  or  alkaline  solu- 
tions of  zinc  salts  by  addition  either  of  hydrogen  sulphide  or  of  a 
soluble  alkaline  sulphide.*  When  so  precipitated,  it  forms  a  white 
powder  ;  in  a  crystalline  state  it  is  found  as  zinc-blende,  a  substance 
which  has  much  the  same  color  as  ordinary  resin. 

Cadmium  sulphide  is  precipitated  from  solutions  of  cadmium 
salts,,  even  if  the  latter  are  slightly  acid,  for  the  substance  is  insolu- 
ble in  dilute  acids.  The  precipitate  is  yellow  in  color. 

Mercury  forms  two  classes  of  compounds,  mercurous  compounds, 
derived  from  the  metal  in  its  monovalent  state,  and  mercuric  com- 
pounds, derived  from  divalent  mercury.  The  same  distinction,  it 
will  be  remembered,  existed  between  cuprous  and  cupric  derivatives. 

Oxides  of  mercury.  Mercurous  oxide  (  Hg2  0),  mercuric  oxide 
(  Hg  0).  The  former  is  produced  by  adding  potassium  hydroxide 
to  a  solution  of  mercurous  nitrate,  the  hydroxide,  which  would  be 
expected,  at  once  breaking  down  into  water  and  the  oxide  :  — 

1.     2  Hg  N03  +  2  KOH  =  2  KST03  +  2  Hg  OH, 

2. 


Mercurous  oxide  is  a  black  powder  which  is  quite  unstable  ;  when 
exposed  to  the  light,  it  breaks  down  into  mercury  and  mercuric 

oxide:-  Hg20  =  Hg  +  HgO.:f 

*  If  the  zinc  salt  is  the  salt  of  a  strong  acid,  such  as  hydrochloric,  nitric, 
or  sulphuric,  only  a  portion  of  the  zinc  is  precipitated  as  the  sulphide,  by 
means  of  hydrogen  sulphide;  for,  as  will  be  seen  from  the  following  reaction, 
a  portion  of  the  acid  is  set  free  during  the  reaction,  and  the  acid  which  is 
formed  decomposes  the  precipitated  sulphide  in  order  to  form  once  more  a 
soluble  salt  of  zinc  :  — 

Zn  SO4  +  H,,  S  =  Zn  S  +  H2  SO4  , 

Zn  S  +  H2  SO*  =  Zn  SO4  +~H.2  S. 

This  difficulty  is  not  encountered  if  the  zinc  salt  of  a  weak  acid  is  used,  or 
if  ammonia  is  previously  added  so  as  to  neutralize  any  acid  which  may  be 
liberated. 

t  It  will  be  remembered  that  the  same  is  true  of  the  formation  of  the 
oxide  of  silver;  when  a  soluble  hydroxide  is  added  to  the  solution  of  a  silver 
salt,  not  the  hydroxide,  but  the  oxide,  is  precipitated  :  — 

2  Ag  NO3  +  2  KOH  =  2  KNO3  +  Ag2  O  +  H.2  O. 

J  This  change  reminds  us  of  the  ones  \vhich  we  encountered  with  many 
of  the  oxides  and  acids  of  the  not-metals. 


MERCUEOUS   CHLORIDE.  431 

Addition  of  acids  produces  mercurous  salts  ;  oxidizing  agents  change 
mercurous  compounds  into  mercuric  compounds. 

Mercuric  oxide  exists  in  two  forms,  according  to  the  method  of 
its  preparation;  the  one  is  red  and  of  crystalline  structure,  the 
other  is  yellow  and  amorphous.  The  red  oxide  can  be  prepared 
either  by  heating  mercury  to  just  below  its  boiling  point  in  the 
presence  of  oxygen,  when,  after  a  long  time,  it  becomes  covered  with 
crystalline  scales  of  the  substance  sought  ;  or  by  heating  mercuric 
nitrate,  which  salt  breaks  down  into  mercuric  oxide  and  nitrogen 
peroxide  (see  page  201)  :  — 


The  yellow  oxide  of  mercury  is  produced  by  precipitation  from 
solutions  of  mercuric  salts  after  the  addition  of  a  soluble  hydroxide, 
the  hydroxide  at  once  breaking  down  into  the  oxide  and  water  :  — 

1.  Hg(N03)2  +  2  KOH  =  Hg  (OH  )2  +  2  KNO  ,  , 

2.  Hg(OH)2  =  HgO  +  H20. 

Both  varieties  of  mercuric  oxide  turn  black  when  heated  ;  they 
resume  their  usual  color  after  cooling;  at  a  dull  red  heat  they 
decompose  into  mercury  and  oxygen  (see  page  18)  ;  sunlight  has 
the  same  effect  as  heat.  When  mercuric  oxide  is  treated  with  an 
acid  it  produces  mercuric  salts.  All  soluble  mercury  compounds,  as 
well  as  the  metal  itself,  are  extremely  poisonous.  Under  certain 
circumstances  mercuric  oxide  displays  slightly  acidic  properties  ; 
for  instance,  it  is  attacked,  in  small  quantities,  by  fused  potassium 
hydroxide.* 

Chlorides  of  mercury.  Mercurous  chloride  (calomel),  HgCl, 
mercuric  chloride  (corrosive  sublknate),  Hg  C12  . 

Mercurous  chloride  sometimes  occurs  in  a  crystalline  form  as  a 
mineral.  Mercurous  chloride  can  be  prepared  either  by  heating 
mercuric  chloride  with  a  sufficient  quantity  of  mercury  :  f  — 

HgCl2  +  Hg  =  2HgCl, 

or  by  adding  hydrochloric  acid,  or  a  soluble  chloride,  to  a  solution 

*  Precipitated  (yellow)  mercuric  oxide  is  also  slightly  soluble  in  cold 
solutions  of  caustic  alkalies. 

t  This  operation  must  be  carried  on  in  large  flasks  stoppered  with  chalk. 
The  calomel  then  sublimes  from  the  lower  part  of  the  flask  and  collects  on  the 
cooler  portions.  If  the  vessel  is  too  small,  a  portion  of  the  mercurous  salt 
will  evaporate. 


432  MERCUROUS   CHLORIDE. 

containing  a  mercurous  salt,  for  mercurous  chloride  is  insoluble  in 
water  or  dilute  acids  :  — 


H   Cl  =  HgCl-f-HN03, 
Hg  N03  +  Na  Cl  =  Hg  Cl  +  Na  N03  . 

In  this  respect,  then,  mercurous  compounds  are  much  like  those  of 
silver  ;  indeed,  all  of  the  monovalent  heavy  metals  act  alike  in  pro- 
ducing insoluble  chlorides.* 

Mercurous  chloride,  when  heated,  evaporates  without  previously 
melting.  The  specific  gravity  of  the  vapor  is  8.01  ;  the  calculated 
specific  gravity  for  a  gas  composed  of  molecules  of  the  formula 
HgCl  is  8.14;  in  this  respect  mercurous  chloride  differs  from 
cuprous  chloride,  for  the  vapor  density  of  the  latter  compound  leads 
to  a  formula  Cu2  C12  .  If  mercurous  chloride  is  treated  with  a  solu- 
tion of  ammonia,  it  turns  black  and  produces  an  insoluble  compound 
which  has  the  formula  NH2  Hg2  Cl  :  — 


2  Hg  Cl  +  2  NH3  =  NH2  Hg2  Cl  +  NH4  Cl  ; 

this  substance  is  regarded  as  being  ammonium  chloride  in  which 
two  atoms  of  hydrogen  have  been  replaced  by  two  of  mercury  :  f  — 


1ST 


fH 
H 

H  ammonium  chloride,  and 
H 


H 
H 

Hg  mercurous  chloramide. 

Hg 

Cl 


,61 

If  mercurous  chloride  is  boiled  with  hydrochloric  acid  it  is  converted 
into  mercuric  chloride  and  mercury  :  — 

2HgCl=Hg+HgCl2. 

Sulphuric  acid,  hot  and  concentrated,  changes  it  into  a  mixture  of 
mercuric  sulphate  and  mercuric  chloride,  nitric  acid  into  mercuric 
nitrate  and  mercuric  chloride.  Chlorine  readily  converts  calomel 
into  mercuric  chloride. 

Mercuric  chloride.     This  salt  can  be  produced  either  by  heating 

*  Cuprous  chloride,  Cu  Cl,  argentic  chloride,  Ag  Cl,  aurous  chloride,  Au  Cl, 
mercurous  chloride,  Hg  Cl,  and  thallous  chloride,  Tl  Cl,  are  insoluble. 

t  It  will  be  remembered  that  the  same  power  of  replacing  hydrogen  in 
ammonium  chloride  is  found  in  the  case  of  cuprous  chloride ;  the  formula  of 
this  ammonium  compound  is,  however,  Cu  Cl,  NH3  or  NH3  Cu  Cl.  Argentic 
chloride,  Ag  Cl,  also  possesses  the  power  of  absorbing  ammonia. 


MERCURIC    CHLORIDE.  433 

mercury  in  a  current  of  chlorine,  *  by  dissolving  the  metal  in  aqua 
regia,  or  by  dissolving  mercuric  oxide  in  hydrochloric  acid.  The 
corrosive  sublimate  of  commerce  is  usually  prepared  by  heating 
mercuric  sulphate  with  sodium  chloride  in  a  wide-mouthed  retort ; 
the  mercuric  chloride  sublimes  and  collects  on  the  cold  portions  of 
the  vessel,  while  sodium  sulphate  remains  behind  :  — 

Hg  S04  +  2  Na  Cl  =  Hg  C12  +  Na2  S04 . 

Corrosive  sublimate  is  a  transparent,  crystalline  mass  which  is 
soluble  in  water  and  which  crystallizes  from  that  solvent  when 
evaporated ;  it  crystallizes  from  aqueous  solution  in  thin  prisms ; 
one  hundred  parts  of  water  at  0°  dissolve  5.73  parts,  and  at  ordinary 
temperatures  about  7  parts  of  mercuric  chloride.  The  solutions 
gradually  decompose  when  they  are  exposed  to  the  light,  while  an 
insoluble  basic  chloride  of  mercury  is  formed.  No  change  takes 
place  if  they  are  kept  in  the  dark.  Mercuric  chloride  has  a  great 
tendency  to  form  double  salts  with  the  chlorides  of  other  metals ; 
so,  for  instance,  the  compounds  Na  Cl,  Hg  C12  +  3  H2  0  ;  K  Cl,  Hg 
C12  +  H2  0  are  known ;  in  this  respect  mercuric  chloride  has 
strongly  marked  acidic  properties,  although,  when  heated  with  phos- 
phorus pentachloride,  it  can  act  as  a  base,  for  it  produces  the 
double  salt  3  Hg  C12 ,  2  P  C15  .f  Mercuric  chloride  is  so  volatile 
that  a  portion  passes  off  when  its  solution  is  evaporated  with  hy- 
drochloric acid  ;  this  loss  mav  be  prevented  by  adding  potassium  or 
sodium  chloride,  for  the  double  chloride  which  is  formed  is  not  vol- 
atile. Reducing  agents  readily  convert  mercuric  chloride  into  mer- 
curous  chloride.  Stannous  chtoride  changes  corrosive  sublimate 
into  calomel :  — 

2  Hg  C12  +  Sn  C12  =  2  Hg  Cl  -h  Sn  C14 , 

and  if  an  excess  of  the  reagent  is  added,  the  mercurous  chloride  is 
finally  changed  to  mercury  (see  page  314)  :  — 

2  Hg  Cl  +  Sn  C12  =  2  Hg  +  Sn  C14 . 

*  The  mercury  burns  with  a  pale  flame,  and  forms  a  white  sublimate  of 
mercuric  chloride. 

t  3  Hg  C12,  2  P  C15  =  Hg3  (P  C18  )2 .  Compare  this  formula  with  that  of  the 
tertiary  phosphate  ;  viz.,  Hg3  (  PO4  )2. 


434  MERCUROUS   NITRATE. 

Ammonia  produces  a  white  precipitate  of  mercuric  chloramide, 
NH2  Hg  Cl.  This  substance  is  analogous  to  the  corresponding 
mercurous  compound  with  the  exception  that  one  atom  of  bi- 
valent mercury  takes  the  place  of  two  atoms  of  hydrogen.  The 
distinction  between  the  two  is  made  apparent  by  the  following 
formulse  :  — 


HI 

H 

Hg  mercurous  chloramide  and  N-[  TT    mercuric  chloramide* 

Her  MS 

»  PI 

Cl 


The  iodides  of  mercury  are  similar  to  the  chlorides.  Mercurous 
iodide,  formed  by  the  direct  union  of  mercury  and  iodine,  readily 
breaks  down  into  mercury  and  mercuric  iodide  :  — 


(the  chloride  suffers  the  same  change  when  heated).  Mercuric 
iodide  displays  even  a  greater  tendency  to  form  double  salts  (in 
which  it  plays  the  part  of  an  acidic  anhydride)  than  does  mercuric 
chloride.  The  fact  that  the  halides  of  so  many  metals  have  acidic 
properties,  while  the  oxides  do  not,  is  not  difficult  to  comprehend  if 
we  remember  that  chlorine,  bromine,  and  iodine  belong  to  the  most 
not-metallic  family  with  which  we  are  acquainted,  so  that  the  halo- 
gen compounds  should  be  more  negative  than  are  those  of  oxygen. 
Some  compounds  are  known  in  which  the  oxide  of  a  metal  is  the 
base,  and  the  chloride,  bromide,  or  iodide  is  the  acidic  anhydride. 
Such  compounds  are  the  double'  oxychlorides  of  mercury.  An 
example  of  one  of  these  compounds  is  Hg  0,  2  Hg  C12  ,  the  insoluble 
precipitate  formed  by  the  gradual  decomposition  of  a  mercuric 
chloride  solution. 

Mercurous  nitrate  is  produced  by  the  action  of  dilute  nitric  acid 
upon  mercury.  When  diluted  with  water,  it  forms  basic  nitrates 
which  are  soluble  with  difficulty.  These  basic  nitrates  vary  in  com- 

*  A  number  of  similar  compounds  derived  from  ammonia  are  known  ;  for 
example,  dimercuric  ammonium  chloride  (  NHg.2Cl)  and  similar  salts  derived 
from  other  acids  have  been  described,  but  for  these  a  larger  text-book  must  be 
consulted. 


MERCURIC    SALTS.  435 

position  according  to  the  temperature  or  the  amount  of  water  used. 
The  simplest  compound  has  the  formula  HgN03,  Hg2  0  -f-  H20, 
and  may  possibly  be  a  tertiary  salt  of  an  hypothetical  ortho-nitric 
acid,  H3N04  (see  page  205)  :  - 


Mercuric  nitrate  is  produced  by  dissolving  mercury  in  concen- 
trated nitric  acid  ;  when  dissolved  in  an  excess  of  water  it  forms  an 
insoluble  basic  nitrate. 

Mercuric  cyanide,  Hg  (CN  )2  ,  is  of  importance  because  it  is  the 
only  cyanide  of  the  heavy  metals  which  is  soluble  in  water.  For 
this  reason  it  is  a  very  useful  reagent  in  chemical  analysis.  The 
cyanide  can  be  readily  produced  by  dissolving  mercuric  oxide  in  a 
solution  of  hydrocyanic  acid.  When  heated,  it  decomposes  into 
mercury  and  cyanogen  (see  page  294).  Mercuric  cyanide  readily 
forms  double  salts  with  the  cyanides  of  other  metals. 

Mercurous  sulphide  has  not,  as  yet,  been  prepared.  Hydrogen 
sulphide,  when  passed  through  a  solution  of  a  mercurous  salt,  pre- 
cipitates a  mixture  of  mercuric  sulphide  and  mercury.  Mercuric 
sulphide  is  the  chief  ore  of  mercury  ;  it  is  a  red,  crystalline  mineral, 
termed  cinnabar.  When  hydrogen  sulphide  is  added  to  the  slightly 
acid  solution  of  a  mercuric  salt,  mercuric  sulphide  is  precipitated  in 
the  form  of  a  black  powder  ;  the  latter  changes  into  the  red  variety 
by  heating,  or  by  treating  for  a  long  time  with  solutions  of  caustic 
alkalies.  Exposure  to  the  light  gradually  changes  the  red  sulphide 
into  the  black  variety.  Black  mercuric  sulphide  is  also  produced 
by  rubbing  mercury  and  sulphur  together.  Mercuric  sulphide  is 
not  attacked  by  dilute  acids  ;  concentrated  nitric  acid  in  part  dis- 
solves it,  and  in  part  converts  it  into  a  white,  insoluble  double  com- 
pound of  mercuric  nitrate  and  mercuric  sulphide;  this  compound 
has  the  formula  Hg  (  N  03)2  2  Hg  S.  Cinnabar,  when  heated,  turns 
black,  and,  unless  the  temperature  was  too  high,  resumes  its  original 
red  color  on  cooling. 

All  mercury  compounds,  if  they  are  salts  of  volatile  acids,  are 
volatile  ;  if,  on  the  other  hand,  they  are  salts  of  not-volatile  acids, 
either  the  acids  themselves  or  their  decomposition  products  remain 
after  heating  (see  page  190). 

*  See  Remsen  ;  Chemistry,  p.  627. 


436 


ZINC;    CADMIUM;    MERC UK Y ;    TABLE    OF. 


TYPICAL    COMPOUNDS    IN    THIS    FAMILY. 


OXIDES. 

HYDBOXIDES. 

CHLOBIDEB. 

SULPHATES. 

SULPHIDES. 

Zinc 
Cadmium 
Mercury 

ZnO 
CdO 
HgO* 

Zn  (OH)2t 
Cd  (OH)2 

Zn  C12  J 
Cd  C12  J 
HgCl2§ 

Zn  SO4  +  7  H2  O 
3  Cd  S04  +  8  H2  O 
HgS04 

ZnS** 
CdS 
HgS 

*  Exists  in  two  varieties,  yellow  and  red. 

t  Both  basic  and  acidic  in  its  character:  Zn  (OH)2  +  2HC1  =  ZnCl2  + 
2  H2  O  and  Zn  (OH)2  +  2  KOH  =  Zn  (OK)2  +  2  H2  O. 

J  Readily  form  double  chlorides  with  ammonium  chloride.  The  latter  are 
not  decomposed  by  ammonia. 

§  Forms  mercuric  chloramide  (NH2HgCl)  with  ammonia. 

**  Soluble  in  dilute  acids. 

All  of  the  salts  of  the  elements  belonging  to  this  group  show  a  great  ten- 
dency to  produce  double  salts. 

MEKCUROUS  COMPOUNDS  (not  typical). 

Hg2  O,  mercurous  oxide. 

Hg  Cl,  mercurous  chloride. 

HgNO3 ,  mercurous  nitrate. 

Mercurous  chloride  is  insoluble  in  water.  When  covered  with  ammonia 
solution  it  forms  mercurous  chloramide  (NH2  Hg2Cl),  the  nitrate  forms  basic 
salts  on  addition  of  water. 


THE   GADOLLNTTE   EARTHS.  437 


CHAPTER   LVI. 

THE   ELEMENTS   BELONGING  TO  THE  PRIMARY  GROUPS  OP  THE 
FAMILIES   HE.,   IV.,   AND  V.,   OF   THE   LONG   PERIODS. 

THE  elements  comprising  the  primary  group  of  the  third  family 
bear  somewhat  the  same  relation  to  boron  and  aluminium  that  cal- 
cium, strontium,  and  barium  do  to  beryllium  and  magnesium.  The 
elements  are :  — 

Scandium ;     symbol,  Sc ;     atomic  weight,    44  ; 
Yttrium;        symbol,  Y;     atomic  weight,    89.1; 
Lanthanum ;  symbol,  La  ;  atomic  weight,  138.2  ; 
Ytterbium;     symbol,  Yb;  atomic  weight,  173. 

All  of  these  elements  are  extremely  rare;  they  occur  in  an 
ortho-silicate  known  as  gadolinite,  the  most  common  formula  of 
which  is  Be2  (  Y0)2  Fe  (Si  04)2 ,  and  also  in  an  extremely  complicated 
salt  of  titanic  acid  known  as  euxenite.  Considerable  uncertainty 
has  existed  as  to  the  number  of  elements  really  contained  in  this 
and  in  the  following  group ;  for  some  investigators,  by  reason  of 
the  peculiarities  of  the  absorption  spectra  of  some  of  the  salts 
of  these  metals,  have  undertaken  to  prove  that  the  usually  accepted 
number  must  be  largely  increased.1* 

Scandium,  it  will  be  remembered,  was  one  of  the  elements  pre- 
dicted by  Mendelejeff  (see  page  373).  The  typical  oxide  of  these 
elements  is  M2  03 ,  corresponding  to  B2  03  and  A12  O3 ;  this  oxide  is 
basic  in  its  character ;  it  does  not  dissolve  in  caustic  alkalies  ;  so  that 
these  elements  are  more  metallic  than  is  aluminium.  The  hydrox- 
ides have  the  formula  M  (OH)3,  the  sulphates,  M2  (S04)3,  and,  un- 
like the  sulphates  of  the  secondary  group  of  this  family  ;  t  they  do 

*  Elements  occurring  in  gadolinite  and  belonging  to  this  group  are  gado- 
linium, 156.1;  samarium,  150;  terbium,  160;  erbium,  166.3.  For  the  descrip- 
tion of  these  elements  the  original  literature  must  be  consulted.  For  methods 
of  separation,  see  Kriiss ;  Liebig's  Annalen ;  265, 1.  Didymium  also  was  among 
the  elements  of  the  gadolinite  group;  this  substance  has  recently  been  sepa- 
rated into  two  elements,  neodymium  and  praseodymium. 

t  Sulphates  of  aluminium,  gallium,  indium,  thallium. 


438  ELEMENTS   OF  TITANIUM  GKOUP. 

not  form  alums ;  this  distinction  is  similar  to  that  existing  between 
the  sulphates  of  calcium,  strontium,  and  barium,  and  those  of  zinc, 
cadmium,  and  mercury;  for  the  sulphates  of  the  first  three  are 
insoluble,  while  those  of  the  second  three  are  soluble,  and  have 
a  great  tendency  to  form  double  salts  with  the  sulphates  of  the 
alkalies. 

The  elements  comprising  the  primary  group  of  family  IV.  are :  — 

Titanium ;     symbol,  Ti ;    atomic  weight,    48  ; 
Zirconium ;  symbol,  Zr ;    atomic  weight,    90.6 ; 
Cerium ;        symbol,  Ce  ;    atomic  weight,  140.2 ; 
Thorium ;      symbol,  Th ;  atomic  weight,  232.6. 

Of  the  compounds  of  these  four  elements,  those  of  titanium  are 
by  far  the  most  common ;  indeed,  compounds  containing  titanium 
are  not  at  all  infrequent,  for  the  element  occurs  in  many  iron  ores ; 
titanic  iron  ore  is  looked  upon  as  being  ferrous  titanate,  Fe  Ti  03 ; 
the  compound  is,  however,  isomorphous  with  ferric  oxide,  and  fre- 
quently occurs  in  company  with  that  extremely  important  sub- 
stance ;  furthermore,  titanic  oxide  is  often  a  constituent  of  magnetic 
iron  ore,  Fe3  O4 ,  which  latter  substance  is  a  member  of  the  spinell 
group  (see  page  339),  so  that  it  seems  not  improbable  that  titanic 
iron  is  really  derived  from  a  hydroxide,  Ti  0  (OH  ),  which  is  analo- 
gous to  Al  0  (OH )  ;  if  this  relationship  is  granted,  then  titanium 
can,  under  certain  circumstances,  act  as  a  trivalent  element,  and 
this  behavior  would  bring  it  in  line  with  cerium  (see  below). 
Titanium  dioxide,  Ti  02 ,  is  found  in  two  mineral  forms,  brookite 
and  anatas,*  and  as  a  polymeric  form,  Ti2  04 ,  which  is  called  rutile, 
and  which  is  isomorphous  with  an  ortho-silicate  of  zirconium, 
Zr  Si  O4 ,  known  as  zircon,  and  with  tinstone  (see  page  312).  The 
relationship  between  these  three  oxides,  all  of  which  belong  to  ele- 
ments in  the  same  family  and  are  isomorphous,  becomes  apparent 
if  we  double  the  formula  of  the  oxide  of  tin,  thus :  — 


SnSn04,  tinstone, 
Ti  Ti  04 ,  rutile, 
Zr  Si  04 ,  zircon. 


*  These  two  oxides  of  titanium  are  not  isomorphous  with  quartz  and 
tridymite,  yet  the  form  of  brookite  is  so  close  to  that  of  tridymite  that  isomor- 
phism is  considered  possible  (Groth).  (See  page  304.) 


TITANIUM   GROUP  ;   COMPOUNDS   OF. 


439 


Cerium  occurs  in  gadolinite  and  also  in  a  silicate  termed  cerite. 

The  compounds  of  the  elements  of  this  group  are  analogous  to 
those  of  silicon ;  this  connection  will  be  seen  from  the  following 
table :  — 


OXIDES. 

CHLORIDES. 

FLUORIDES. 

HYDROXIDES. 

METAHYDBOXIDES. 

Silicon 
Titanium 

Si02 
TiO2 

7r  O 

SiCl4* 
TiCl4* 

7r  Cl    t 

SiF4§ 
TiCl4§ 
Zr  PI    8 

Si(OH)4 
Ti(OH)4 
Zr  (OH)4 

Si03H2 

Ti08H2(?) 

Th  O 

Th  Cl    t 

Th  CL  S 

Th  (OH)4 

*  Decomposed  by  water. 

t  Partially  decomposed  by  water;  forming  a  basic  chloride: 


(  Cl  +  HOH 
-I   Cl  +  HOH 

Zr  )   Cl 

lei 


OH 

"1   C1 

lei 


Zr 


OH 

ATT 

X;  i 

Cl 


=zr 


Cl 


t  Not  decomposed  except  by  hot  water. 

§  All  of  these  fluorides  behave  exactly  as  does  silicon  tetrafluoride ;  they 
form  compounds  analogous  to  fluosilicic  acid  and  the  fluosilicates  (see  page 
303):  — 

H2  Si  F6 ,  fluosilicic  acid ;      K2  Si  F6 ,    potassium  fluosilicate. 
H2  Ti  F6 ,  fluotitanic  acid ;    K2  Ti  F6  ,   potassium  fluotitanate. 
K2  Zr  F6  ,    potassium  fluozirconate. 
K2  Th  F6 ,  potassium  fluothorate. 

The  compounds  of  cerium  are  omitted  from  this  table  because  the  chemistry 
of  this  element  is  not  yet  sufficiently  clear  for  purposes  of  comparison. 

The  above  table  and  the  explanatory  notes  very  plainly  show 
the  intimate  family  connection  between  silicon  and  the  three  fol- 
lowing elements ;  the  oxides  of  titanium  and  zirconium  resemble 
silicon  dioxide  in  the  fact  that,  after  they  have  been  heated,  they 
are  insoluble  in  water  and  even  in  the  strongest  acids  or  alkalies ;  to 
be  brought  in  solution  they  must  be  heated  with  concentrated  sul- 
phuric acid  for  a  long  time,  or  they  must  be  fused  with  alkalies  ;  the 
oxide  of  thorium  is  somewhat  less  obstinate.  The  last  three  ele- 
ments in  the  above  table,  having  higher  atomic  weights  than  silicon, 
are  also  more  metallic  in  their  nature ;  their  hydroxides  are,  conse- 
quently, both  weak  acids  and  weak  bases.  The  salts  in  which  they 
act  as  acids  are  but  little  known ;  indeed,  it  is  doubtful  if  thorium 
hydroxide  has  acidic  properties.  The  salts  of  the  alkalies,  so  far 
as  known,  correspond  to  the  metasilicates,  and  hence  have  the 


440  ELEMENTS    OF   VANADIUM   GROUP. 

general  formula  M2  X03 .  Among  the  salts  derived  from  the  hy- 
droxides, acting  as  bases,  the  sulphates,  with  the  general  formula 
M(S04)2,  are  perhaps  the  most  prominent.  The  chemistry  of 
cerium  is,  as  yet,  uncertain  in  many  respects ;  the  element  forms 
two  series  of  compounds,  in  one  of  which,  presenting  compounds 
such  as  Ce203,  CeCl3,  Ce(N03)3,  it  is  trivalent,  and  resembles 
lanthanum,  an  element  in  the  preceding  family ;  in  the  other  it  is 
tetravalent,  and  by  means  of  compounds  Ce  F4 ,  Ce  02 ,  Ce  (  S04 )  2 ,  it 
falls  in  line  with  the  family  numbered  IV.  It  may  be  added  that 
titanium  likewise  forms  more  than  one  oxide ;  for  a  compound, 
Ti203,  and  a  sulphate,  Ti2  (S04)  3,  have  been  described. 

The  elements  comprising  the  primary  group  of  family  V.  are  :  — 

Vanadium ;  atomic  weight,     51.4  ;  symbol,  V. 

Columbium ;         atomic  weight,     94;  symbol,  Cb. 

Neodymium ;        atomic  weight,  140.5 ;  symbol,  Nd. 

Praseodymium  ;  atomic  weight,  143.5  ;  symbol,  Pr. 

Tantalum,  atomic  weight,  182.6  ;  symbol,  Ta. 

Although  more  metallic  in  their  nature  than  the  elements  form- 
ing the  secondary  group  of  this  family,*  the  four  elements,  with 
perhaps  the  exception  of  neodymium  and  praseodymium,  bear  many 
points  of  resemblance  to  this  group.  In  some  respects  the  two  lat- 
ter substances  are  very  much  like  cerium  and  lanthanum,  being 
trivalent  in  most  of  their  compounds.  Vanadium  is  the  best  known 
of  all  of  these  elements.  This  element  is  as  much  like  arsenic  or 
antimony  as  titanium  is  like  silicon ;  in  very  many  respects  it  is,  in- 
deed, like  the  typical  element  of  the  family,  nitrogen,  for  it  forms 
as  many  oxides  as  the  latter,  and  these  oxides  have  similar  for- 
mulae, thus : — 

OXIDES    OF   NITROGEN.  OXIDES    OF   VANADIUM. 

N2  0  V2  0 

NO(N202)  V202 

N203  V203 

N02,(N204)  V204 

N205  V205 

Vanadium  occurs  in  nature  chiefly  as  the  lead,  zinc,  or  bismuth 
salt  of  vanadic  acid.  Vanadinite  is  a  double  salt  composed  of  lead 

*  Arsenic,  antimony,  and  bismuth. 


VANADIC   ACIDS. 


441 


vanadate  and  lead  chloride,  having  the  formula  3  Pb3(V04)2-j- 
PbCl2*.  The  element  forms  three  chlorides,  V  C12,  V  C13,  and 
VC14.  The  most  important  compounds  of  vanadium  are  the  va- 
nadic  acids,  which  correspond  to  those  of  phosphorus  and  of  arsenic, 
the  latter  being  elements  belonging  to  the  same  family :  - 


OXIDES. 

META-ACIDS. 

OETHO-ACIDS. 

PYKO-ACIDS. 

Phosphorus 
Arsenic 
Vanadium 

P205 

As205 
V205 

HPO3 
HAsO3 
HV03 

H3P04 

H3  As  O4 
H3V04 

H4P207 
H4  As.,  O7 
H4V207 

Vanadates  of  all  these  acids  are  known ;  for  instance,  we  have  :  — 

Na  V03 ,  sodium  metavanadate ; 
Na3  V04 ,  sodium  orthovanadate ; 
Na4  V2  07  sodium  pyrovanadate. 

In  addition  to  the  above,  however,  more  complicated  salts  of 
polyvanadic  acids,  which  are  formed  in  the  same  manner  as  the 
polysilicic  acids,  are  known.  Free  metavanadic  acid  is  a  golden 
yellow,  crystalline  compound.  Reducing  agents  readily  change 
vanadic  acid  into  the  lower  acid,  V2  04 . 

Columbium  and  tantalum  are  very  rare  elements  which  differ 
from  vanadium  just  as  much  as  antimony  does  from  arsenic;  for 
they  are  able  to  form  pentachlorides  and  pentabromides,  while  the 
pentahalides  of  vanadium  have  not  as  yet  been  prepared.  Colum- 
bium is  also  frequently  termed  niobium.  It  occurs  in  the  min- 
eral columbite,  which  is  a  metacolumbate  of  iron,  having  the 
formula  Fe  (Cb  03)2;  tantalum  is  found  as  a  metatantalate  of  iron, 
Fe(Ta03)2. 

The  brief  mention  of  the  very  rare  metals  which  have  been 
discussed  in  this  chapter  is  sufficient  to  demonstrate  the  family 
relationship  existing  between  them  and  the  much  more  common 
elements  which  were  described  in  the  first  portion  of  the  book  ;  of 
course,  they  form  a  great  number  of  compounds,  some  of  them  very 
complicated,  which  cannot  be  taken  up  in  a  book  of  this  kind ;  for 
this  study  a  large  manual  must  be  consulted. 

*  Yanadinite  is  isomorphous  with  apatite,  which  has  a  formula  3  Ca3 
(  PO4  )2  +  Ca  C12 ,  calcium  replacing  lead,  and  vanadium  replacing  phos~ 
phorus  isomorphously. 


442  GADOLINITE   METALS   AXD   PERIODIC    SYSTEM. 

The  uncertainty  which  exists  as  to  the  number  of  elements 
which  can  be  isolated  from  gadolinite,  euxenite,  and  allied  minerals 
leaves  that  portion  of  the  periodic  system  in  which  these  elements 
find  their  places  in  a  very  chaotic  condition.  Certainly,  if  all  of 
these  supposed  elements  should  finally  prove  themselves  to  be  such, 
we  should  then  have  to  conclude  that  several  elements,  with  very 
nearly  the  same  atomic  weight,  must  occupy  the  same  position  in 
the  periodic  system.  This  discovery  need  not  overthrow  existing 
relations  with  well-known  elements,  for  the  following  are  facts 
which  cannot  be  controverted. 

The  periodic  system,  as  it  now  stands,  is  undoubtedly  a  natural 
grouping  of  the  elements.  The  fact  that  certain  portions  of  the 
system  are  now  completely  filled  with  well-known  elements  does 
not  prove  that  other  elements,  intermediate  to  known  groups,  will 
never  be  discovered.  Indeed,  were  we  acquainted  with  seven  hun- 
dred elements,  instead  of  seventy,  these  seven  hundred  would  no 
more  be  unconnected  individuals  than  are  our  present  number. 
They  too  would  fall  into  an  arrangement,  in  periods  and  families, 
as  the  elements  now  do ;  only,  with  such  a  large  number  of  indi- 
viduals, the  change  from  family  to  family  would  not  present  such 
an  abrupt  transition  as  at  present. 


ELEMENTS   OF   CHROMIUM   GKOUP. 


443 


CHAPTER   LVII. 

THE  ELEMENTS  BELONGING  TO  THE  PRIMARY  GROUP  OP  THE 

VI.  FAMILY. 

Chromium  ;    symbol,  Or  ;    atomic  weight,  52.1 ; 

Molybdenum  ;   symbol,  Mo  ;    atomic  iveight,  96  j 

Tungsten  (  Wolfram)  ;    symbol,  W ;    atomic  weight,  184  ; 

Uranium  ;  symbol,  U  ;  atomic  weight,  239.6. 
THE  typical  elements  belonging  to  the  sixth  family,  in  the  short 
periods,  are  oxygen  and  sulphur  ;  and,  as  has  been  shown  by  the 
arrangement  of  the  periodic  system  given  on  page  363,  the  individ- 
uals more  immediately  connected  with  those  two  elements  are  sele- 
nium and  tellurium,  while  chromium,  molybdenum,  tungsten,  and 
uranium,  having  their  positions  near  the  middle  of  the  long  periods, 
vary  much  more  from  the  character  of  the  types  in  the  short  periods 
than  do  the  metals  which  have  been  discussed  in  the  preceding  chap- 
ters. The  metallic  nature  of  the  elements  forming  the  primary 
group  of  the  sixth  family  is  most  apparent  in  the  behavior  of  the 
lower  oxides  and  in  the  salts  derived  from  these  ;  the  highest  oxide 
of  each  element  is  the  typical  one,  X03 ;  this  compound  displays 
the  character  of  an  acidic  anhydride,  although,  in  the  case  of  the 
most  metallic  element  of  the  family  (uranium),  it  is  also  basic  under 
some  circumstances.  The  salts  derived  from  the  typical  oxide,  in 
formulae,  correspond  to  the  sulphates,  selenates,  and  tellurates,  and 
in  some  instances  to  the  pyrosulphates  (see  page  154);  this  relation- 
ship is  made  apparent  by  the  following  table  :  — 


OXIDES. 

ACIDS. 

SALTS. 

DI-8ALT8. 

OXIDES. 

ACID8.- 

SALTS. 

PI-SALTS. 

S03 

H2S04 

M,  S04 

M2S207 

Cr03 

H2CrO4* 

M2Cr04 

M2  Cr2  07 

(Se03) 

H2Se04 

M3SeO4 

MoO3 

H2MoO4t 

M2  Mo  O4 

M2Mo2O7 

Te03 

H2TeO4 

M2  Te  04 

WO  3 

HaW04t 

M3WO4 

M2W207 

UO3 

M2U2O7 

*  The  acid  is  not  known ;  when  liberated  from  its  salts  it  breaks  down 
into  its  anhydride,  Cr  O3 ,  and  water. 

t  The  ortho-acid,  H4  XO5 ,  is  also  known. 

Chromium,  molybdenum,  and  tungsten  further  form  a  number  of  very 
complicated  salts  derived  from  polyacids  which  are  produced  in  a  manner 
analogous  to  the  polysilicates  (see -page  307). 


444  CHROMIUM  GROUP;  OCCURRENCE. 

The  most  marked  characteristic  of  these  elements,  and  one  they 
share  with  the  others  having  their  position  a»fc  the  middle  of  the  long 
periods,  lies  in  the  power  which  the  individuals  possess  of  forming 
several  series  of  compounds,  in  each  of  which  they  exhibit  a  differ- 
ent valence ;  so,  for  instance,  molybdenum  forms  the  chlorides 
Mo  C12 ,  Mo  C13 ,  Mo  C14 ,  and  Mo  Cl  6 ,  while  tungsten  exhibits  W  C12 , 
WC14,  WC15,  and  WC16;  these  two  elements  possess  chlorides, 
therefore,  in  which  they  are  respectively  quinquivalent  and  hexava- 
lent;  and,  passing  backward  from  these,  we  find  a  series  of  com- 
pounds in  which  the  valence  successively  diminishes  by  one  until  a 
minimum  is  reached  at  divalence.  It  will  be  remembered  that  sul- 
phur shows  some  resemblance  to  molybdenum  and  tungsten  by 
forming  three  chlorides  of  the  formulae  S2  C12 ,  S  C12 ,  and  S  C14  re- 
spectively, but  the  latter  is  capable  of  existence  only  at  very  low 
temperatures;  it  is  not  inconceivable  that,  were  the  proper  condi- 
tions attainable,  a  penta-  and  hexa-chloride  of  sulphur  might  also  be 
produced  (see  page  370). 

None  of  the  elements  of  this  group  occur  uncombined  as  natural 
minerals ;  the  principal  compounds  which  are  found  are  given  in  the 
following  table :  — 

Chromium.  Found  as  chromite  (chromic  iron)  isomorphous  with  spinell, 
formula  Fe  (CrO2)2;*  when  the  chromium  is  replaced  in  part  by  aluminium 
and  by  ferric  iron,  and  the  ferrous  iron  by  magnesium,  the  mineral  is  called 
chromspinell;  chromite  forms  veins  or  imbedded  masses  in  serpentine  rock. 

Crocoite.     Lead  chromate,  Pb  Cr  04  ,t  is  sometimes  found. 

Molybdenum.  Found  as  the  sulphide  MoS2,  called  molybdenite;  as  the 
molybdate  of  lead,  Pb  Mo  O4 ,  called  wulfenite;J  and  as  molybdite,  Mo  O3 , 
the  anhydride  of  molybdic  acid. 

*  This  compound  is  ferrous  chromite,  derived  from  a  hydroxide  of  the 
formula  CrO(OH),  analogous  to  AIO(OH).  The  ferrous  iron  in  chromite 
can  be  replaced  isomorphously  by  divalent  chromium,  the  trivalent  chromium 
in  CrO  (OH)  by  ferric  iron  (see  page  339). 

t  The  chromates  of  the  alkali  metals  are,  without  exception,  isomorphous 
with  the  corresponding  sulphates ;  naturally  occurring  lead  chromate  is,  how- 
ever, not  isomorphous  with  anglesite  (PbS04),  but  artificially  prepared  crys- 
tals of  lead  chromate  have  proven  to  be  isomorphous  'with  the  latter.  The 
isomorphism  of  the  chromates  and  sulphates  clearly  demonstrates  the  family 
connection  between  chromium  and  sulphur. 

|  Wulfenite  is  not  isomorphous  with  crocoite,  but  it  is  sometimes  found 
with  a  contents  of  chromium  replacing  molybdenum.  It  is  isomorphous  with 
the  corresponding  salt  of  tungstic  acid. 


CHROMIUM  GEOUP;   METALLURGY.  445 

Tungsten.  As  scheelite,  Ca  WO4 ;  reinite,  Fe  WO4 ;  and  stoltzite,  Pb  WO4 . 
All  of  these  minerals  are  isomorphous  with  wulf enite.  Tungsten  is  also  found 
as  tungstite,  W03 ,  the  anhydride  of  tungstic  acid. 

Uranium.  As  pitchblende,  (UO2,Pb)  W209;  and  as  the  sulphate  of  ura- 
nium, which  exists  as  an  impure  mineral  sometimes  called  uranocher. 

The  elements  under  discussion  are  not  of  any  commercial  impor- 
tance when  isolated  from  their  compounds.  Chromium  can  be  pre- 
pared by  electrolyzing  the  fused  chloride,  CrCl3,  in  a  manner 
analogous  to  the  preparation  of  the  alkali  metals,  the  alkaline 
earths,  and  of  aluminium ;  or  the  metal  can  be  obtained  by  heating 
the  chloride  with  sodium  or  with  zinc  in  the  absence  of  the  air,  the 
process  being  like  that  formerly  used  for  obtaining  aluminium  (see 
page  333).  Sodium  amalgam,  when  treated  with  chromic  chloride, 
forms  sodium, chloride  and  liberates  chromium,  which  latter  element 
then  forms  an  amalgam  with  mercury.  It  can  be  separated  from 
this  by  distilling  the  mercury  in  a  current  of  hydrogen.  Tungsten, 
molybdenum,  and  uranium  can  be  prepared  by  reducing  the  oxides 
of  these  metals  by  means  of  hydrogen  at  red  heat,  the  elements  in 
question  being  much  more  easily  separated  from  their  oxides  than  is 
chromium  from  its  corresponding  compounds.  In  the  case  of  ura- 
nium, the  element  can  even  be  obtained  by  heating  its  oxides  with 
charcoal.  The  most  important  physical  properties  of  these  elements 
are  given  in  the  following  table  :  — 

Chromium,  specific  gravity  6.8,  crystalline,  of  metallic  appearance,  in- 
fusible. 

Molybdenum,  specific  gravity  8.6,  silver  white,  infusible. 

Tungsten,  specific  gravity  18.1,  steel-colored  plates,  fusible  at  a  high  tem- 
perature. 

Uranium,  specific  gravity  18.4,  white,  metallic  lustre,  fusible  at  a  high 
temperature. 

Chromium,  molybdenum,  tungsten,  and  uranium  all  have  small 
atomic  volumes.  As  was  mentioned  on  page  367,  they  have  their 
places  on  the  descending  branches  and  hear  the  minimum  of  the 
curves  formed  by  using  the  atomic  volumes  as  ordinates,  and  the 
atomic  weights  as  abscissae ;  they  are  therefore  infusible,  or  at  least 
fusible  with  difficulty,  and  they  form  colored  salts. 

Chromium  is  slowly  oxidized  when  heated  in  the  air,  more 
rapidly  in  a  current  of  oxygen ;  the  oxide  which  is  formed  has  the 
formula  Cr203.  The  metal  is  dissolved  by  hydrochloric  acid,  or  by 
hot  sulphuric  acid,  the  chloride,  CrCl3,  or  the  sulphate,  Cr2  (S04)3, 


446  CHROMIC    OXIDE  ;    HYDROXIDE. 

being  produced,  according  to  the  acid  used.  Potassium  nitrate  and 
potassium  chlorate,  when  fused  with  chromium,  give  up  their  oxy- 
gen and  form  potassium  chrornate. 

Molybdenum  is  slowly  attacked  when  heated  in  the  air.  It  is 
readily  converted  into  the  trioxide,  Mo  03 ,  by  oxygen  at  a  high 
temperature.  Nitric  acid  or  aqua  regia  attacks  the  metal  to  form 
molybdic  acid.  Tungsten  behaves  as  does  molybdenum,  while 
uranium  even  dissolves  in  dilute  hydrochloric  or  sulphuric  acid, 
hydrogen  being  at  the  same  time  evolved. 

The  most  important  compounds  of  chromium  are  derived  from 
two  oxides,  chromic  oxide,  Cr2  03,  which  is  mainly  basic  in  its  char- 
acter, and  chromium  trioxide.,  Cr03,  which  acts  as  an  acidic  anhy- 
dride, and  which  is  the  oxide  typical  of  the  family.  In  addition  to 
these  twro,  there  exists  a  chromous  hydroxide,  Cr  (01?  )2  (the  oxide 
corresponding  to  which  is  not  known),  and  a  chromous-chrornic  oxide 
of  the  formula  Cr304. 

Chromic  oxide,  Cr2  03,  is  a  dark  green  powder  which  is  insoluble 
in  acids  when  it  has  been  heated  to  a  high  temperature  (see  page 
338) ;  after  it  has  been  subjected  to  such  treatment,  it  can  be 
brought  into  solution  by  fusion  with  caustic  alkalies,  or  with  the 
primary  sulphate  of  potassium.*  Chromic  oxide  dissolves  in  fused 
glass,  and  imparts  a  fine  green  color  to  the  substance  ;  for  this  rea- 
son it  is  used  as  a  green  paint  for  tinting  porcelain. 

Chromic  hydroxide  can  be  precipitated  from  solutions  of  chromic 
salts  by  the  addition  of  ammonia  water.  When  dry,  it  has  the  com- 
position represented  by  the  formula  Cr  (  OH  )  3  +  2  H2  0.  When 
the  latter  is  heated  to  200°,  it  changes  to  metachromic  hydroxide, 
CrO(OH).  This  substance  corresponds  to  the  similar  aluminium 
compound  (see  page  339).  Chromic  oxide  and  hydroxide  are  both 
basic  and  acidic  in  their  character.  The  freshly  precipitated  hy- 
droxide is  readily  dissolved  by  caustic  alkalies,  forming  deep  green 
solutions  of  the  chromites'  of  the  respective  metals  ;  if  the  solutions 
so  formed  are  allowed  to  stand,  or  if  they  are  boiled,  the  hydroxide 
is  once  more  precipitated,  f  A  number  of  chromites  occur  as  min- 

*  In  this  process  potassium-chrome  alum,  K2  SO4,  Cr2(SO4)  3  +  24  H2  O,  is 
produced. 

t  Differing  from  aluminium  hydroxide,  the  solution  of  which  in  alka- 
lies is  not  decomposed  by  boiling.  When  chromic  oxide  is  present,  alkalies 
are  also  able  to  dissolve  considerable  quantities  of  ferric  oxide. 


CHROMIC    SALTS.  447 

erals ;  the  latter  are  derived  from  metachromic  hydroxide  and  are 
isomorphous  with  spinell ;  these  compounds  can  also  be  artificially 
prepared  by  fusi-ng  chromic  oxide  with  the  metallic  oxide,  which  is 
to  be  used  as  the  base,  boron  trioxide,  B2  03,  being  used  as  a  flux.* 
The  chromic  salts,  in  which  chromic  oxide  acts  as  a  base,  are  pro- 
duced by  dissolving  the  oxide  or  hydroxide  in  acids.  The  most 
important  of  these  are  described  below  :  — 

CHROMIC  CHLORIDE,  Cr  Cl3,t  is  produced  by  burning  chromium  in  an  atmos- 
phere of  chlorine,  or  by  heating  an  intimate  mixture  of  chromic  oxide 
and  charcoal  in  a  current  of  dry  chlorine.  The  salt  so  produced  sub- 
limes in  the  form  of  pink  plates ;  the  chloride  which  is  formed  by  dis- 
solving the  hydroxide  in  water  crystallizes  in  dark  green  needles  of  the 
formula  Cr  C13  +  6  H2  O ;  anhydrous  chromic  chloride  cannot  be  pre- 
pared from  this,  as,  upon  heating,  the  salt  loses  hydrochloric  acid  and 
changes  into  a  basic  chromic  chloride.  The  dry  chloride  is  nearly  in- 
soluble in  water  or  in  acids;  when  heated  in  air,  it  gives  off  chlorine 
and  leaves  chromic  oxide.  The  vapor  density  of  gaseous  chromic 
chloride  shows  that  the  substance  has  a  molecule  corresponding  to 
the  formula  Cr  C13 . 

CHROMIC  SULPHATE,  Cr2  (SO4)3  +  15  H2  O,  is  formed  by  dissolving  chromic 
hydroxide  in  concentrated  sulphuric  acid,  and  then  allowing  the  solu- 
tion to  absorb  moisture  from  the  air;  the  salt  is  reddish  violet  in  color; 
when  heated  to  100°  it  loses  water  of  crystallization,  and  changes  to  a 
green  salt  having  the  composition  Cr2  (SO4)3  +  5  H2  O.  When  a  solu- 
tion of  chromic  sulphate  is  mixed  with  a  solution  of  an  alkaline  sul- 
phate and  evaporated,  an  alum  is  formed  in  which  chromium  has 
taken  the  place  of  aluminium ;  an  example  of  such  a  salt  is  K2  SO4 , 
Cr2  (SO4)3  +  24  H2  O  (see  page  339). 

When  solutions  of  caustic  alkalies  are  added  to  solutions  of 
chromic  salts,  a  precipitate  of  chromic  hydroxide  is  produced ;  the 
latter  is  soluble  in  an  excess  of  the  precipitating  medium ;  on  the 
other  hand,  chromic  hydroxide  does  not  dissolve  in  ammonia  solu- 
tion, and  can,  as  a  consequence,  be  precipitated  from  solutions  of 
chromic  salts  by  the  addition  of  that  reagent,  even  in  excess;  t  al- 

*  Zinc  chromite,  Zn(CrO2)o,  and  manganous  chromite,  Mn(CrO2)2, 
have  been  prepared  in  this  way. 

t  The  formula  Cr2  C16  was  formerly  assigned  to  this  compound ;  but  the 
latest  determinations  of  the  specific  gravity  of  this  substance,  while  in  the 
state  of  a  vapor,  show  it  to  be  Cr  C13  (see  page  335).  Nilsson  and  Pettersson; 
Comptes  Rendus;  107,  529. 

J  The  precipitation  of  chromic  hydroxide  is  very  much  retarded,  and  may 
be  entirely  prevented,  by  the  presence  of  not-volatile  organic  acids  such  as 
citric  acid,  tartaric  acid,  oxalic  acid,  etc. 


448  CHROMIC    ACID. 

kaline  sulphide  solutions  or  solutions  of  the  carbonates  precipitate 
chromic  hydroxide  for  reasons  identical  with  those  mentioned  in 
the  chapter  on  aluminium  (see  page  341).*  When  in  alkaline  solu- 
tion, chromic  hydroxide  is  completely  oxidized  to  a  chromate  by  the 
addition  of  chlorine  or  bromine  ;  the  same  change  can  also  be  brought 
about  by  the  addition  of  other  oxidizing  agents,  f  or  by  fusion  with 
potassium  nitrate  or  chlorate. 

Chromic  acid,  H2Cr04,  is  not  known;  its  anhydride,  Cr03, 
chromium  trioxide,  is  produced  when  the  acid  is  liberated  from  its 
salts ;  this  is  best  accomplished  by  adding  tolerably  concentrated 
sulphuric  acid  to  a  solution  of  potassium  or  sodium  dichromate :  — 

K2  Cr2  07  +  H2  S04  =  K2  S04  +  H2  0  +  5  Cr  03 . 

The  anhydride  crystallizes  in  beautiful  carmine-red  needles  which 
melt  at  193°,  forming  a  dark  red  fluid  which  loses  oxygen  at  250°, 
and  changes  to  green  chromic  oxide ;  the  oxide  is  readily  soluble  in 
water,  the  solubilfty  being  diminished  by  the  addition  of  sulphuric 
acid.J 

Chromium  trioxide  is  a  most  energetic  oxidizing  agent ;    even 
dilute   solutions   instantly  change   sulphurous   acid  into  sulphuric 
acid:—  L     2Cr03  +  3H2S03  =  Cr203         +3HaS04; 
2.       Cr2  03  +  3  H2  S04  =  O2  (S04 ),  +  3  H2  0. 

In  the  same  way  hydrogen  sulphide  is  oxidized,  while  sulphur  is 
separated.  Many  organic  substances  are  also  readily  attacked  by 
chromium  trioxide,  §  that  substance  being  at  the  same  time  reduced 
to  chromic  oxide  (respectively  to  the  chromic  salts  which  would  be 
formed  by  the  acids  which  may  be  present). 

Although  chromic  acid  is  unknown,  an  acid  chloride  of  chromium 
called  chromyl  chloride,  Cr  02  C12 ,  which  may  be  considered  as  being 
chromic  acid  in  which  the  two  hydroxyl  groups  have  been  replaced 

*  Chromium  is  completely  separated  from  solutions  of  chromic  salts  by 
the  addition  of  freshly  precipitated  barium  carbonate.  The  chromium  sepa- 
rates as  chromic  hydroxide  mixed  with  basic  chromic  salt. 

t  For  instance,  lead  superoxide,  Pb  O.2 ,  when  lead  chromate  is  produced. 

J  Chromium  trioxide  is  least  soluble  in  sulphuric  acid  of  about  85  per  cent. 

§  For  instance,  alcohol  is  oxidized  to  aldehyde  and  to  acetic  acid ;  the  re- 
lationship between  these  three  compounds  can  be  seen  from  the  following 
structural  formulae:  — 

CH3  -  CH2  OH  +  O  =  CH3  -  COH  +  H2  O.     CH3  COH  +  O  =  CH3  COOH. 
Alcohol,  Aldehyde,  Acetic  acid. 


CHROMYL   CHLORIDE  ;    CHROMATES.  449 

by  chlorine,  can  be  produced  by  adding  concentrated  sulphuric  acid 
to  an  intimate  mixture  of  sodium  chloride  with  potassium  dichro- 
mate. Chromyl  chloride  is  a  dark  red  fluid  which  boils  at  118°,  and 
which  instantly  decomposes  into  chromium  trioxide  and  hydrochloric 
acid  on  the  addition  of  water  :  — 

Cr  02  C12  +  H2  0  =  2  H  Cl  +  Cr  03 . 
In  structure,  this  compound  is  analogous  to  sulphuryl  chloride :  — 

0 
Cl  I  Cl 

Chromyl  chloride          and          Sulphuryl  chloride  (see  page  157.) 
Salts  of  another  acid  chloride  of  chromium,  in  which  only  one 
hydroxyl  group  in  a  formula  weight  has  been  replaced  by  chlorine, 
are  also  described.     These  salts  correspond  to  those  of  chlor-sul- 
phonic  acid  (see  page  157). 

The  chromates  are  derived  from  a  hypothetical  dibasic  acid 
analogous  in  formula  to  sulphuric  acid :  — 

fOH  fOH 

s  Jo 

S  |0 

tOH  lOH 

Sulphuric  acid        and  Chromic  acid; 

and  the  dichromates  form  a  dichromic  acid,  H2  Cr2  07  (also  hypo- 
thetical), which  is  analogous  to  disulphuric  acid,  H2  S2  O7  (see  page 
154).*  The  most  important  chromates  and  dichromates  are  those 
of  potassium  and  of  sodium. 

POTASSIUM  CHROMATE  is  a  yellow,  crystalline  salt,  which  is  readily  soluble 
in  water,  and  which  is  isomorphous  with  potassium  sulphate.  Upon 
addition  of  dilute  acids  it  is  converted  into  the  dichromate :  — 

2  K2  Cr  O4  +  2  HNO3  =  K2  Cr2  O7  +  2  KNO3  +  H2  O.t 

*  Tri-  and  poly-chromates  have  been  described  (see  page  307).  A  di-sul- 
phuric  acid  in  which  a  portion  of  the  sulphur  has  been  replaced  by  chromium 
is  also  known. 

t  In  this  reaction  the  primary  chromate  of  potassium,  KHCr  O4,  may  be 
considered  to  be  the  first  product :  — 

K2  Cr  O4  +  HNO3  =  KHCr  O4  +  KNO3 . 

Two  formula  weights  of  this  primary  salt  would  then  separate  water,  leaving 
the  dichromate  :  -  2  KHCr  ^  =  ^  ^  ^  +  ^  Q 


450  BICHROMATES  ;    LEAD   CHIIOMATE. 

On  the  other  hand,  potassium  dichromate  is  changed  to  the  chromate 
by  alkalies,  thus :  K2  Cr.2  O7  +  2  KOH  =  2  K2  Cr  O4  +  H2  O. 
POTASSIUM  DICHROMATE  crystallizes  in  red  prisms  or  plates  which  are 
soluble  in  water.  It  melts  at  a  low  red  heat  without  decomposition, 
and  it  loses  oxygen  and  is  converted  into  a  mixture  of  chromic  oxide 
and  potassium  chromate  at  a  bright  white  heat :  — 

2  K2  Cr2  O7  =  2  K2  Cr  O4  +  Cr2  O3  +  3  O. 

It  also  liberates  oxygen  when  it  is  heated  with  non-oxidizable  mineral 
acids  (such  as  sulphuric  acid).  With  sulphuric  acid,  chromic  sulphate 
(respectively  potassium-chrome  alum)  is  formed:  — 

K2  Cr2O7  +  4  H2  SO  4=  K2  SO4 ,  Cr2  (SO4  )8  +  4  H2  O  +  3  O.* 

Of  course,  an  acid  solution  of  potassium  dichromate  can  oxidize  the 
hydrogen  compounds  of  the  not-metals;  these  reactions  are  discussed 
on  pages  60  and  97.  Potassium  dichromate  is  extensively  used  for  the 
preparation  of  battery  fluids ;  however,  of  late  years,  the  more  soluble 
sodium  dichromate  is  taking  the  place  of  the  potassium  salt. 

The  reactions  shown  by  sodium  dichromate  are  identical  with 
those  of  potassium  dichromate. 

THE  CHROMATE  OF  LEAD  is  insoluble  in  water,  and  is  produced  by  adding 
a  soluble  chromate  or  dichromate  to  the  solution  of  a  lead  salt,  as 
follows :  — 

1.  K2  Cr2  O4  +  Pb  (NO3)2  =  2  K  NO3  +  Pb  Cr  O4 . 

2.  K2  Cr.2  O7  +  2  Pb  (NO3)2  +  H2  O  =  2  KXO3  +  2  HNO3  +  2  Pb  Cr  O4 . 

In  the  second  reaction,  therefore,  free  acid  is  produced;  the  same  is 
true  in  other  cases  where,  by  double  decomposition,  an  insoluble 

*  In  this  reaction  the  oxygen  present  in  the  chromate,  in  excess  of  that 
necessary  to  form  chromic  oxide,  passes  off.  The  same  is  true  of  the  other 
reactions  in  which  potassium  dichromate,  or  dichromates  in  general,  are  used 
as  oxidizing  agents.  One  formula  weight  of  potassium  dichromate,  therefore, 
has  three  atoms  of  oxygen  at  its  disposal  for  oxidizing  purposes.  In  formulat- 
ing reactions  this  fact  is  the  essential  one  to  be  taken  into  consideration.  One 
formula  weight  of  potassium  dichromate  will  therefore  oxidize  three  of  sul- 
phurous acid  to  sulphuric  acid,  or  three  molecules  of  alcohol  to  aldehyde,  etc. 
Concentrated  hydrochloric  acid  is  oxidized  to  chlorine  by  the  dichromate;  in 
the  latter  case,  of  course,  six  molecules  of  hydrochloric  acid  are  changed  to 
chlorine  by  the  dichromate,  for:  — 

6HC1  +  30=3H20+  6C1; 

however,  the  excess  of  hydrochloric  acid  will  subsequently  form  potassium 
chloride  and  chromic  chloride  with  the  bases  present,  so  that  the  complete 
reaction  would  be  represented  as  follows :  — 

K2  Cr2  O7  +  14  HC1  =  2  KC1  +  2  Cr  C13  +-G  Cl  +  7  H2  O. 


CHROMOUS   COMPOUNDS.  451 

cliromate  can  be  formed,  for  the  precipitation  takes  place  by  the 
addition  of  either  a  soluble  cliromate  or  dichromate  to  the  solution 
containing  the  salt  of  the  metal  capable  of  forming  such  an  insoluble 
cliromate.  Lead  cliromate  possesses  a  bright  yellow  color,  which 
makes  it  useful  as  a  paint  (chrome  yellow*);  addition  of  potassium 
or  sodium  hydroxide  to  chrome  yellow  changes  that  substance  into 
an  insoluble  basic  lead  chroniate  which  is  termed  chrome  red  :  — 

2  Pb  Cr  O4  +  2  KOH  =  (Pb  OH)2  Cr  O4  +  K2  Cr  O4  . 

BARIUM  CHKOMATE,  Ba  Cr  O4  ,  is  insoluble  in  water,  but  soluble  in  hydro- 
chloric or  nitric  acid.  In  this  way  the  salt  differs  from  the  equally 
insoluble  barium  sulphate,  for  the  latter  is  insoluble  both  in  water 
and  in  acids.  The  solubility  of  barium  cliromate  is  not  increased  by 
the  presence  of  other  salts  in  solution  in  the  supernatent  fluid,  but  the 
solubility  of  the  sulphate  is  increased. 

The  chromous  compounds,  which,  are  derived  from  chromous 
hydroxide  (CrOH)2,  are  of  far  less  importance  than  either  the 
chromic  salts  or  the  chromates. 

CHROMOUS  CHLORIDE,  CrCl2,  is  produced  when  metallic  chromium  is 
heated  in  a  current  of  dry  hydrochloric  acid,  or  when  chromic  chlo- 
ride, Cr  C13  ,  is  reduced  by  means  of  a  current  of  dry  hydrogen.  The 
substance  produces  a  blue  solution  in  water ;  the  latter,  however, 
rapidly  turns  green,  owing  to  the  absorption  of  oxygen  and  the  forma- 
tion of  a  basic  chromic  chloride;  addition  of  alkaline  hydroxides  to 
this  solution  precipitates  brownish-yellow  chromous  hydroxide.  One 
of  the  chief  characteristics  of  all  chromous  compounds  is  the  extreme 
ease  with  which  they  take  up  oxygen  in  order  to  produce  chromic 
salts.  A  number  of  chromous  salts  of  acids  other  than  hydrochloric 
acid  have  also  been  prepared.  The  specific  gravity  of  the  vapor 
of  chromous  chloride  shows  that  it  has  a  molecule  corresponding  to 
the  formula  CrCl2.  (ISTilsson  and  Pettersson;  Comptes  Rend.;  107, 
529.) 

An  oxide  of  chromium  having  the  formula  Cr3  04  is  also  known. 
This  substance  is  regarded  as  being  composed  of  chromous  and 
chromic  oxides  :  -  Cr  0  +  Cr2  03  =  O3  04 ; 

so  that,  in  this  compound,  chromous  oxide  plays  the  part  of  a  base, 
and  chromic  oxide  that  of  an  acidic  anhydride. 

The  compounds  of  chromium  which  are  used  in  the  arts  are  pre- 
pared from  chromic  iron.  The  latter  substance  is  finely  ground, 
washed  with  water,  and  then  intimately  mixed  with  potash-lime ;  * 

*  A  mixture  of  potassium  and  calcium  hydroxides  produced  by  slaking 
quick-lime  with  a  potassium  hydroxide  solution;  soda-lime  is  produced  by 
using  sodium  instead  of  potassium  hydroxide. 


452  MOLYBDENUM;   COMPOUNDS   OF. 

it  is  afterward  dried  at  150°,  and  finally  heated  in  reverberatory 
furnaces,  the  oxygen  so  supplied  changing  the  chromic  compound 
into  potassium  chroniate  and  calcium  chromate ;  *  after  heating  for 
a  sufficient  length  of  time,  the  mass  is  extracted  with  the  least 
quantity  of  boiling  water,  and  the  solution  so  produced  is  then 
treated  with  potassium  sulphate.  By  this  means  the  calcium 
chromate  is  converted  into  potassium  chromate  ;  the  latter  salt  is 
finally  changed  into  potassium  dichromate  by  the  addition  of  sul- 
phuric acid.  Potassium  and  sodium  dichromates  are  used  for  the 
preparation  of  various  paints  (chrome  yellow,  chrome  red,  etc.) ;  in 
a  number  of  processes  of  dyeing;  in  the  preparation  of  chromic 
oxide,  which  is  used  for  porcelain  painting ;  as  oxidizing  agents ; 
and  in  a  number  of  other  ways. 

The  compounds  of  molybdenum,  tungsten,  and  uranium  are  not, 
by  any  means,  so  important  as  are  those  of  chromium. 

Molybdenum  forms  the  following  oxides  :  — 

1.  Molybdenum  monoxide,  MoO;  brown  (nearly  black)  in  color. 

2.  Molybdic  oxide,  Mo.2  O3 ;  black  in  color. 

3.  Molybdenum  dioxide,  MoO2;  dark  brown  in  color,  t 

4.  Molybdenum  trioxide,  Mo  O$ ;  white  in  color. 

The  last  of  these  is  produced  by  roasting  finely  powdered  molyb- 
denite, by  which  means  the  sulphur  is  burned  off,  and  impure  yellow 
molybdenum  trioxide  is  left ;  the  latter  is  extracted  with  ammonia 
water,  which  produces  ammonium  molybdate.  The  ammonium 
molybdate  is  recrystallized,  and  is  then  converted  into  molybdic  acid 
by  gently  heating  it :  — 

(  NH4  )2  Mo  04  =  H2  Mo  04  +  2  NH3 . 

The  trioxide  is  produced  by  dehydrating  the  latter  compound.  The 
oxide  is  difficultly  soluble  in  water,  is  white  and  of  crystalline 
structure  ;  reducing  agents  change  it  to  the  second  oxide,  Mo2  03 ;  and 
this,  when  gently  heated,  takes  up  oxygen  and  forms  the  dioxide, 
Mo  02 ;  the  first  oxide  is  formed  by  adding  a  concentrated  solution 
of  potassium  hydroxide  to  molybdenum  dichloride,  Mo  C12 .  Hydrox- 
ides, Mo  (OH  )8  and  Mo  (OH  )4 ,  corresponding  to  Mo2  03  and  M02 
are  also  known. 

*  Of  course,  in  the  preparation  of  sodium  dichromate,  soda-lime  is  used, 
t  A  blue  oxide,  Mo3  O8 ,  is  also  described. 


MOLYBDIC    ACIDS.  453 

Molybdenum  forms  the  following  chlorides  :  — 

Mo  C15  ,  produced  by  passing  dry  chlorine  over  heated  molybdenum. 
Mo  CU ,  produced  by  heating  the  trichloride  in  a  current  of  carbon  dioxide.* 
Mo  C\s  ,  produced  by  heating  the  pentachloride  in  a  current  of  hydrogen 
at  250°. 

Mo  C12 ,  produced  by  heating  the  trichloride  in  a  current  of  carbon  dioxide.* 

Molybdenum  pentachloride  boils  at  268°,  changing  into  a  dark 
brown  gas  which  has  a  specific  gravity  of  9.4  at  350° ;  this  vapor 
density  corresponds  to  a  molecule  having  the  formula  Mo  C15 ,  so 
that  molybdenum  affords  an  example  of  an  element  belonging  to  the 
sulphur  family  which  enters  into  the  formation  of  a  compound  in 
which  the  individual  is  pentavalent  (see  page  370). 

Unlike  chromium  trioxide,  the  trioxide  of  molybdenum  is  capa- 
ble of  forming  a  number  of  hydrated  acids  which,  in  formula,  corre- 
spond to  the  hydrated  sulphuric  acids ;  this  relationship  is  made 
plain  by  the  following  table  :  — 

MOLYBDIC    ACIDS.  SULPHURIC   ACIDS  (see  page  145). 

Mo  O3       4-  H2  O  =  H2  Mo  O4 .  SO3       +  H2  O  =  H2  SO4 . 

H2  MoO4  +  H2O  =  H4  Mo  O5.  H2  SO4  +  H2  O  =  H4  SO6  . 

H4Mo05  +  H~  O  =  H6Mo  O6 .  H4  SO5  +  H2  O  =  H6SO6  . 

The  acid  having  the  formula  H4Mo05  is  formed  by  adding 
nitric  acid  to  a  solution  of  sodium  or  potassium  molybdate.  It  is 
nearly  insoluble  in  water,  and  when  dried  in  vacua  loses  water  while 
changing  to  H2  Mo  04 .  H6  Mo  06  is  soluble  in  water,  and  is  pro- 
duced by  dialysis  of  a  solution  of  molybdic  acid  in  a  manner  similar 
to  the  separation  of  soluble  silicic  acid  (see  page  305).  Salts  of  a 
number  of  complicated  molybdic  acids  are  also  known ;  the  formulae 
of  a  few  of  these  are  given  in  the  following  table :  — 

Na2  Mo2  O7 ,   sodium  dimolybdate.t 
Na2  Mo3O10 ,  sodium  trimolybdate. 
]STa2  Mo4  O18 ,  sodium  tetramolybdate. 
]STa2Mo8O25,  sodium  octomolybdate. 

The  method  of  formation  of  these  polymolybdates  is  similar  to  that 
of  the  polysilicates. 

Acid  solutions  of  molybdic  acid  are  readily  reduced  by  means  of 
metallic  zinc  or  tin.  During  such  a  reduction  the  color  of  the  solu- 

*  2  Mo  C13  =  Mo  C12  +  Mo  C14 .      These  chlorides  must  be  produced  in  the 
absence  of  free  oxygen ;  otherwise  the  oxy-chlorides  are  formed. 
t  Corresponding  to  the  disulphate  and  dichromate. 


454  MOLYBDIC   ACIDS. 

tion  at  first  becomes  blue,  then  green,  and  finally  black,  at  which 
stage  of  the  reaction  the  monoxide  Mo  0  is  precipitated. 

When  nitric  acid  is  added  to  a  solution  of  ammonium  molybdate, 
ammonium  tetramolybdate  [(NH4)2  Mo4013]  is  produced.  The 
latter  substance  is  of  great  importance  in  analytical  chemistry,  for 
the  reason  that,  upon  addition  of  phosphoric  acid  or  a  soluble  phos- 
phate, the  phosphoric  acid  is  completely  separated  as  a  constituent 
of  a  yellow  precipitate  known  as  ammonium  phospho-molybdate. 
The  latter  substance  has  the  composition  expressed  by  the  formula 
(  NH4  )3  P04 , 11  Mo  03 .  *  When  this  salt  is  treated  with  aqua  regia 
the  free  phospho-niolybdic  acid,  H3P04,  11  Mo03,  goes  into  solu- 
tion. The  power  which  molybdic  acid  possesses  of  uniting  with 
other  acids  to  form  complicated  compounds  is  most  important  in 
the  case  of  phosphoric  acid ;  however,  it  is  not  confined  to  that 
substance  alone,  for  similar  unions  of  molybdic  acid  with  arsenic 
and  silicic  acids  are  also  known. 

Molybdenum  forms  three  sulphides,  a  disulphide,  Mo  S2,  a 
trisulphide,  MoS3,  and  a  tetrasulphide,  MoS4.  If,  in  the  latter 
compound,  we  regard  sulphur  as  having  a  valence  of  two,  then 
molybdenum  may  possibly  be  octovalent.  In  speculating  as  regards 
the  valence  of  molybdenum  in  such  a  compound,  we  must  always 
bear  in  mind,  however,  that,  as  it  cannot  be  vaporized,  its  molecular 
weight  is  unknown. 

The  tetrasulphide  of  molybdenum  can  act  as  an  acidic  anhydride, 
for  it  forms  a  potassium  salt  of  the  formula  K2  Mo  S5 . 

Tungsten  produces  only  two  oxides,  a  dioxide,  W02 ,  and  a 
trioxide,  W03 ,  f  the  latter  being  the  anhydride  of  tungstic  acid  ; 
the  element,  however,  is  able  to  enter  into  four  chlorides,  namely  :  — 

Tungsten  dichloride,  W  C12 . 
Tungsten  tetrachloride,  W  C14 . 
Tungsten  pentachloride,  W  C15 . 
Tungsten  hexachloride,  W  C16 . 

*  More  complicated  formulae  have  recently  been  assigned  to  ammonium 
phospho-molybdate  and  the  phospho-molybdic  acid.  They  are  (  NH4  )8  PO4 , 
12  Mo  O3  +  (  NH4)2  HPO4 ,  12  Mo  O3  +  8  H2  O  and  (HN4)2  HPO4 , 12  Mo  O3  + 
29  H2  O. 

t  An  oxide,  W3  O8 ,  corresponding  to  Mo3  O8 ,  is  known.  This  oxide  is 
probably  a  combination  of  W  O2  acting  as  a  base  and  W  O3  acting  as  an  anhy- 
dride:— 


TUNGSTIC   ACIDS. 


455 


The  last  compound  is  produced  by  heating  tungsten  in  a  cur- 
rent of  chlorine ;  it  boils  at  346°,  and  has  a  vapor  density  which 
corresponds  to  the  molecular  weight  expressed  by  the  formula 
W  C16 ;  so  that  tungsten  must  be  hexavalent  in  the  hexachloride. 

Provided  we  regard  the  atoms  of  oxygen  as  always  being  diva- 
lent, then  the  highest  valence  of  the  elements  of  the  sulphur  group, 
when  in  combination  with  oxygen,  is  six ;  and  in  one  instance  at 
least,  as  is  shown  by  the  existence  of  W  C16 ,  the  valence  toward 
chlorine  also  reaches  that  number.  It  seems  not  improbable,  there- 
fore, that,  were  the  proper  conditions  attainable,  the  remaining  ele- 
ments of  this  family  would  also  be  able  to  produce  compounds 
which,  in  each  molecule,  would  contain  six  atoms  of  chlorine.  We 
should  then  have  a  series  of  chlorides,  derived  from  members  of  the 
first  six  families,  which  would  exactly  correspond  to  the  oxides,  two 
chlorine  atoms  taking  the  place  of  one  of  oxygen.  This  will  be 
made  clear  from  the  following  general  formulae :  — 


FAMILY. 

1. 

2. 

3. 

4. 

5. 

6. 

Chlorides 

RC1 

RC12 

RC13 

RC14 

RC15 

RC^ 

Oxides 

R,0 

RO 

R203 

RO2 

R205 

R03 

The  oxides  of  the  seventh  family,  K2  07  ,  have  as  yet  no  corre- 
sponding halide,  but  it  seems  not  impossible  that  some  of  the 
missing  compounds  will  ultimately  be  discovered.  In  the  first  six 
families,  however,  the  highest  valence  of  the  elements  toward  oxy- 
gen and  toward  chlorine  is  given  by  the  number  of  the  family  to  which 
respectively  each  group  of  elements  belongs. 

Two  tungstic  acids,  which  correspond  to  H2  S04  and  H4  S05  , 
are  known  ;  they  are  H2  W04  and  H4  W05  .  The  first  of  these  is 
a  yellow  powder,  which  is  produced  by  decomposing  the  aqueous 
solution  of  an  alkaline  tungstate  with  an  excess  of  hot  acid  :  — 


Na2  W04 


H2  W04  . 


The  second  is  produced  by  using  cold  instead  of  hot  acid.  A  num- 
ber of  polytungstates  which  are  similar  to  the  polymolybdates  are 
also  known.  Tungstic  acid  has  the  same  ability  of  uniting  with 
other  acids  to  form  complicated  compounds  as  is  possessed  by  molyb- 


456  URANIUM;    OXIDES    OF. 

die  acid.  We  are  acquainted,  for  instance,  with  phosphotungstic 
acid,  arsenotungstic  acid,  silicotungstic  acid,  etc.  Of  these  com- 
pounds, perhaps  the  most  important  is  silicotungstic  acid,  the  sodium 
salt  of  which  is  formed  by  boiling  a  polytungstate  of  sodium  *  with 
precipitated  silicic  acid ;  the  latter  dissolving  to  form  sodium  silico- 
tungstate,  having  a  formula  Na8  Si  Wi2  042  +  22  H2  0.  This  salt 
is  extremely  soluble  in  water,  and  its  solution  has  a  high  specific 
gravity. f  The  free  acid  is  soluble  in  ether. 

Uranium  has  the  highest  atomic  weight,  and  hence  the  most 
metallic  nature  of  any  of  the  elements  under  discussion.  As  a  con- 
sequence, its  trioxide  can  act  both  as  a  base  and  as  an  acid. 

The  oxides  of  uranium  correspond  exactly  to  those  of  tungsten. 
They  are  U02 ,  uranous  oxide,  U03 ,  uranic  oxide,  and  U3  08 ,  which 
is  considered  to  be  uranous-uranic  oxide.  Only  three  chlorides  of 
uranium,  namely,  a  trichloride,  U  C13 ,  a  tetrachloride,  U  C14 ,  and  a 
pentachloride,  U  C15 ,  are  known. 

Uranous  oxide,  U  02 ,  is  basic  in  its  properties  and  forms  salts  of 
the  general  formula  U  X4 ,  where  X  represents  the  remainder  of  a 
monobasic  acid  after  the  removal  of  hydrogen.  The  uranous  salts 
are  colored  green,  and  are  easily  oxidized  to  compounds  derived  from 
uranic  oxide. 

When  uranic  oxide  enters  into  combination  with  acids,  it  forms 
basic  salts  in  which  the  divalent  radicle  =U02  plays  the  part  of  a 
divalent  metal.  This  radicle  is  called  uranyl,  and  its  relationship 
to  its  salts  is  similar  to  that  of  the  univalent  radicle  stibionyl, 
SbO  — ,  which  was  described  on  page  253.  The  resemblance 
between  the  radicle  uranyl  and  the  atoms  of  divalent  metals  can 
be  seen  by  comparing  the  following  formulae  of  uranyl  and  calcium 
salts :  — 

(U02)  (N03)2 ,  uranyl  nitrate.  Ca  (NO3)2 ,  calcium  nitrate. 

(U02)  S04 ,  uranyl  sulphate.  Ca  S04 ,  calcium  sulphate. 

(U02)3  (P04)2 ,  uranyl  phosphate.          Ca3  (P04)2 ,  calcium  phosphate. 

*  Sodium  paratungstate,  Na2  W12  O41 .  This  salt  is  formed  hy  fusing  to- 
gether reinite  (FeWO4  )  and  sodium  carbonate.  It  finds  extensive  applica- 
tion in  the  manufacture  of  a  fireproof  sizing  for  inflammable  materials. 

t  The  specific  gravity  is  3.  when  the  solution  is  saturated  at  ordinary  tem- 
peratures. 


URANYL   HYDROXIDE  ;   URANATES.  457 

The  hydroxide  from  which  the  uranyl  salts  are  derived  can  be  com- 
pared to  calcium  hydroxide,  the  divalent  group  =U02,  taking  the 
place  of  one  atom  of  calcium  :  — 

U02(OH)2       and         Ca(OH)2 

Uranyl  hydroxide      and    calcium  hydroxide. 

Uranyl  hydroxide  can,  therefore,  dissolve  in  acids  (for  example, 
nitric  acid),  and  it  then  forms  uranyl  nitrate,  exactly  as  calcium 
hydroxide  can  dissolve  in  the  same  reagent  to  form  calcium 
nitrate  ;  the  two  reactions  may  consequently  be  expressed  as  fol- 
lows :  — 

U02  (OH)2  +  2  HM)3  =  U02  (N03)2  +  2  H2  0  ; 
Ca     (OH)2  +  2HN03  =  Ca     (N03)2  +  2H20. 


The  uranyl  salts  are  yellow,  with  a  green  fluorescence. 

Uranyl  hydroxide,  in  addition  to  being  a  base,  is,  however,  also 
an  acid  ;  it  dissolves  in  strong  bases  to  form  uranates  ;  and,  by  writ- 
ing the  formula  in  a  manner  slightly  different  from  that  given  above, 
it  will  be  seen  that  uranyl  hydroxide  is  also  uranic  acid,  and  there- 
fore it  corresponds  to  sulphuric  acid  :  — 

H2  S04  ,  sulphuric  acid,  and  H2  U04  ,  uranic  acid. 

The  uranates,  however,  in  formula  resemble  the  disulphates  and 
dichromates,  and  not  the  sulphates  and  chromates.  When  a  uranyl 
salt  is  treated  with  a  solution  of  a  caustic  alkali,  the  first  change 
would  be  the  formation  of  uranyl  hydroxide  (uranic  acid)  :  — 


U02  (N03)2  +  2  KOH  =  U02  (OH)2 
this,  however,  reacts  with  the  alkali  to  form  a  diuranate  :  — 
2  U02  (OH  )2  +  2  KOH  =  K2  U2  07  +  3  H2  0.* 

The  chief  compounds  discussed  in  the  last  chapter  are  given  in 
the  following  table  :  — 

*  Sodium  uranate,  when  fused  with  glass,  imparts  a  yellow  tint  with  green 
fluorescence  to  the  same. 


458  ELEMENTS   OF   CHKOMIUM   FAMILY;    TABLE   OF. 


CHLOBIDES. 

OXIDES. 

CrCLj 
CrCl3 

MoCl2 
MoCl3 
MoCl4 
MoCl6 

WC12 

(CfO>» 

Cr203f 

MoO 

Mo2  O3  * 
MoO2§ 

WC14 

WC15 
WC16 

UC14 
UC15 

WO2 

U02* 

Cr03t 

Mo03t 

W03t 

U03J 

Cr304 

Mo308 

W308 

U308 

The  oxides  on  the  last  line  are  considered  to  be  combinations  of  two  other 
oxides. 


ACIDS. 

DI-  AND  POLY-ACIDS. 

CrO8 

Mo08 

W03 

TJ03 

All  of  these  elements  also  form  salts 

H2MoO4 
H4  Mo  O5 

H2W04 
H4W05 

H2  U04tt 
H4  U05** 

derived  from  a  di-acid  having  the  gen- 
eral formula  H3  X2  O7  ,  and  they  also 
form  salts  derived    from  complicated 



H6Mo06** 



poly-acids;    the  latter  are  formed  by 

• 

uniting  3,  4,  5,  etc.,  formula  weights  of 

the  acids  H2  XO4  ,  and  then  separating 

water  until  a  dibasic  acid  is  left. 

The  above  acids  are  all  dibasic  (see  page  140). 

*  Basic.  t  Acidic.  J  Basic  and  acidic. 

§  The  solution  of  Mo  O2  reddens  litmus  paper,  but  possesses  no  other  acid 
properties. 

**  Existence  doubtful. 

tt  Also  acts  as  a  basic  hydroxide,  uranyl  hydroxide,  U02  (OH)2. 


MANGANESE  ;   OCCURRENCE.  459 


CHAPTER   LVIIL 

THE   ELEMENT   FORMING   THE   PRIMARY  GROUP    OF   THE 
SEVENTH   FAMILY. 

I 

Manganese  ;  symbol,  Mn ;  atomic  weight,  55. 

ONLY  one  element  which  should  undoubtedly  have  its  place  in 
the  primary  group  of  the  seventh  family  has,  as  yet,  been  discov- 
ered ;  and  that  element  is  manganese.  Manganese,  having  its  place 
at  the  middle  of  one  of  the  long  periods,  must  necessarily  differ 
very  markedly  from  the  typical  elements  of  the  family  (i.e.,  from  the 
halogens)  ;  and,  indeed,  a  great  variation  from  the  properties  of  the 
latter  is  to  be  expected  even  without  any  such  consideration,  for 
manganese  is  metallic  in  its  nature,  while  fluorine,  chlorine,  bromine, 
and  iodine  are  the  most  negative  of  all  elements.  As  a  conse- 
quence, we  should  expect  the  greatest  resemblance  between  the  halo- 
gens and  manganese  to  lie  in  the  derivatives  of  the  highest  oxides. 
In  these  compounds,  the  metallic  nature  of  manganese  is  almost 
entirely  overshadowed  by  the  negative  elements  with  which  it  is 
combined,  so  that  the  permanganates,  R  Mn  04  ,*  in  many  respects 
(such  as  isomorphism,  solubility,  etc.),  are  very  much  like  the  per- 
chlorates.  The  lower  oxides  of  manganese,  on  the  other  hand, 
bear  no  resemblance  to  the  halogen  oxides ;  in  fact,  their  nearest 
prototypes  are  to  be  found  among  the  oxides  of  iron,  chromium, 
cobalt,  nickel,  or  lead,  while  in  many  respects  Mn  0  acts  very  much 
like  the  oxides  of  calcium,  magnesium,  or  zinc.  The  typical  oxide 
of  the  seventh  family,  therefore,  is  X2  07 ;  in  no  case,  excepting 
that  of  manganese,  has  it  been  isolated ;  it  is,  however,  known  in 
its  derivatives  (permanganates,  perchlorates,  periodates,  and  the 
corresponding  acids). 

Manganese  is  never  found  as  the  uncombined  element.  The 
chief  minerals  in  which  it  occurs  are  given  in  the  following 
table :  - 

*  R  represents  a  univalent  metal. 


460  MANGANESE;  PROPERTIES. 

Braunite,  Mn2  03 .  Pyrolusite  (polianite),  Mn  02 . 

Hausmannite,  Mn304.  Manganite,  MiiO  (OH).* 

Pyrolusite  and  inanganite  are  the  most  important  ores  of  man- 
ganese. They  both  occur  in  large  beds  and  in  veins.  Manganous 
oxide,  Mn  0,  is  also  sometimes  found  as  a  mineral  termed  manga- 
nosite.  The  carbonate,  rhodochrosite,  Mn  C03,  belongs  to  the  calcite 
group  (see  page  416),  while  manganocalcite,  (Mn,  Ca)  C03,  probably 
is  isomorphous  with  arragonite. 

The  element  itself  is  very  difficult  to  separate  from  its  ores ;  for 
such  powerful  reducing  agents  as  red-hot  charcoal  or  hydrogen  are 
able  to  change  the  higher  oxides  into  manganous  oxide  only,  but 
not  into  manganese.!  Manganese  can,  however,  be  isolated  either 
by  heating  manganous  chloride  with  sodium,  or  by  electrolysis  of 
the  fused  chloride  or  fluoride. 

Manganese  is  a  grayish-white  metal,  which  somewhat  resembles 
cast  iron ;  it  is  crystalline  in  structure  and  brittle,  although  it  pos- 
sesses a  certain  amount  of  toughness.  The  specific  gravity  is  about 
8,  and  its  atomic  value  6.9.  Manganese  is,  therefore,  at  the  mini- 
mum of  one  of  the  curves  of  atomic  volumes  ;  the  element  with  next 
smaller  atomic  weight  has  a  larger  atomic  volume,  and,  as  a  conse- 
quence, manganese  is  difficult  to  fuse,  and  forms  colored  salts.  The 
melting  point  of  manganese  lies  at  about  1900°  ;  this  is  probably 
somewhat  higher  than  that  of  iron.  Pure  manganese,  after  polishing, 
rapidly  becomes  dull  when  exposed  to  the  air,  owing  to  oxidation. 
The  metal  is  energetically  attacked  by  acids.  $  Pure  manganese 
has  no  technical  application ;  an  alloy  of  manganese  and  iron 
(f  erro-manganese,  spiegeleisen)  is,  however,  of  the  greatest  commercial 
importance  for  the  manufacture  of  Bessemer  steel. 

Manganese  forms  the  following  oxides  :  — 

Mn  O,  manganous  oxide. 
Mn2  O3 ,  manganic  oxide. 
Mn3  O4 ,  manganous-manganic  oxide. 

MnO2,  manganese  dioxide  (manganese  hyperoxide,  black  oxide  of  man- 
ganese). 

Mn2  O7 ,  permanganic  anhydride. 

*  Corresponding  to  Al  O  (OH).  Trivalent  manganese  can  replace  alu- 
minium and  chromium  isomorphously  in  the  spinells. 

t  The  conversion  of  the  oxides  into  metallic  manganese  by  means  of  char- 
coal takes  place  only  at  a  high  white  heat. 

}  Impure  manganese  decomposes  even  water  very  readily. 


MANGANOUS  COMPOUNDS.  461 

Manganous  oxide  is  basic  in  its  character ;  it  readily  dissolves  in 
acids  to  form  the  manganous  salts ;  this  oxide  can  be  produced 
through  reduction  of  any  one  of  the  higher  oxides  by  heating 
in  a  current  of  hydrogen.  It  is  green,  or  grayish-green,  in  color, 
and  when,  exposed  to  the  air,  it  readily  absorbs  oxygen  to  form 
Mn3  04  .*  Manganous  hydroxide,  Mn  (OH  )2 ,  is  separated  as  a 
white  precipitate  when  alkaline  solutions  or  ammonia  water  are 
added  to  a  solution  containing  a  manganous  salt ;  precipitation  b}^ 
means  of  ammonia  is,  however,  entirely  prevented  by  the  presence 
of  ammonium  salts,  for  manganous  salts  have  a  tendency  to  form 
double  salts  with  the  compounds  of  ammonium,  identical  with  that 
displayed  by  the  similar  compounds  of  zinc  or  of  magnesium  (see 
pages  419  and  438).  When  exposed  to  the  air,  manganous  hy- 
droxide rapidly  turns  brown,  because  it  absorbs  oxygen  and  is  con- 
verted into  manganous-manganic  oxide,  Mn3  04 ;  when  dissolved  in 
acids,  the  corresponding  manganous  salts  are  produced.  The  latter 
do  not.  spontaneously  oxidize  when  exposed  to  the  air. 

MANGANOUS  CHLORIDE  is  contained  in  the  colorless  solutions  obtained  by 
dissolving  any  one  of  the  oxides  of  manganese  in  hydrochloric  acid.t 
When  slowly  evaporated  the  solutions  deposit  pinkish  colored  tablets 
of  the  formula  Mn  C12  +  4  H2  O.  The  anhydrous  salt  cannot  be  ob- 
tained from  these  by  heating,  because  the  chloride,  at  a  temperature 
high  enough  to  drive  off  water  of  crystallization,  loses  chlorine,  ab- 
sorbs oxygen,  and  in  part  changes  into  Mn3  O4 .  In  order  to  obtain 
the  chloride  in  an  anhydrous  condition,  the  water  of  crystallization 
must  be  driven  off  in  a  current  of  dry  hydrochloric  acid  gas.  Man- 
ganous chloride  is  extremely  soluble  in  water,  and  shows  the  greatest 
tendency  to  form  double  chlorides  with  the  corresponding  salts  of 
other  metals. 

MANGANOUS  SULPHATE  is  produced  by  dissolving  any  one  of  the  oxides  of 
manganese  in  hot  sulphuric  acid,t  or  better,  by  dissolving  the  carbon- 
ate in  the  diluted  acid.  Upon  evaporating  and  cooling  to  below  6°, 
crystals  having  the  formula  Mn  SO4  +  7  H2  O  separate.  The  latter 
are  isomorphous  with  the  vitriols  (see  page  417).  If  the  temperature 
of  crystallization  is  between  7°  and  20°,  then  the  crystals  contain  but 

*  This  action  may  become  so  violent  as  to  cause  the  whole  mass  to  glow. 

t  Should  an  oxide  containing  more  oxygen  than  Mn  O  be  dissolved,  the 
excess  of  oxygen  will  oxidize  the  hydrochloric  acid,  liberating  chlorine  (see 
page  58). 

\  When  an  oxide  of  manganese  containing  more  oxygen  than  Mn  O  is 
dissolved,  then  the  excess  of  oxygen  passes  off. 


462  MANGANIC   COMPOUNDS. 

five  molecules  of  water,  and  are  then  isomorphous  with  ordinary  sul- 
phate of  copper  (blue  vitriol;  see  page  403).  The  sulphate  is  readily 
soluble  in  water. 

MANGANOUS  CARBONATE  is  insoluble  in  water,  and  is  therefore  precipi- 
tated from  solutions  of  manganous  salts  by  the  addition  of  a  soluble 
carbonate.  The  naturally  occurring  salt  is  isomorphous  with  calcite. 

MANGANOUS  SULPHIDE  is  insoluble  in  water,  but  soluble  in  dilute  acids  * 
(see  page  100)  ;  it  is  therefore  precipitated  from  solutions  containing 
manganous  salts  by  the  addition  of  an  alkaline  sulphide  solution  :  — 

Mn  SO4  +  (KH4  )2  S  =  (NH4  )2  SO4  +  Mn  S. 

Manganous  sulphide  is  a  flesh-colored  precipitate  which  readily  ab- 
sorbs oxygen  from  the  air,  while,  at  the  same  time,  it  turns  of  a 
brown  color. 

Manganic  oxide,  Mn2  03  ,  occurs  in  nature  as  the  mineral  brau- 
nite,  which  is  the  hardest  ore  of  manganese.  In  the  laboratory  it 
may  be  produced  by  heating  manganous  oxide,  manganese  dioxide, 
or  manganous-manganic  oxide  to  red  heat  in  a  current  of  hydrogen. 
The  oxide  is  black  in  color,  insoluble  in  water,  and,  when  heated  to 
a  white  heat,  changes  to  Mn3  04  .  The  hydroxide  Mn  02  H,  corre- 
sponding to  A102H,  occurs  as  the  mineral  manganite  ;  this  com- 
pound can  also  be  formed  by  slow  oxidation  of  manganous  hydroxide, 
Mn  (OH  )2  ,  in  the  air,  but,  if  the  manganous  hydroxide  is  covered 
with  ammonia  solution,  then  the  product  of  oxidation  is  the  normal 
manganic  hydroxide,  Mn  (OH  )3  .  Both  the  oxide  and  hydroxides 
are  weakly  basic  in  character.  The  salts  derived  from  them  are 
unstable,  and  resemble  those  derived  from  the  oxide  of  aluminium, 
A12  03  .  They  are  decomposed  by  the  addition  of  an  excess  of  water. 
Their  solutions  are  dark  brown  in  color,  and  on  addition  of  alkalies 
precipitate  manganic  hydroxide. 

MANGANIC  CHLORIDE,  Mn  C13  ,  is  produced  by  dissolving  manganic  hydrox- 
ide in  cold  hydrochloric  acid.  The  solution  has  a  dark-brown  color, 
and,  on  standing,  liberates  chlorine,  leaving  manganous  chloride. 

MANGANIC  SULPHATE,  Mn2  (SO4)3,  is  a  dark-green,  amorphous  powder, 
which  is  produced  by  heating  finely  divided  manganese  dioxide  with 
concentrated  sulphuric  acid  to  110°.  A  portion  of  the  oxygen  of  the 
dioxide  then  passes  off,  while  manganese  sulphate  remains.  The  re- 
action may  be  considered  as  taking  place  in  two  stages  :  — 

1.  2MnO2  =  Mn2O3  +  O; 

2.  Mn203 


*  Even  in  acetic  acid  (difference  from  zinc  sulphide). 


MANGANESE   DIOXIDE.  463 

If  the  heating  be  carried  too  far,  more  oxygen  will  be  evolved,  and 
manganous  sulphate,  Mn  SO4  ,  will  remain.  Manganic  sulphate  is  in- 
teresting because  it  forms  compounds  with  the  sulphates  of  the  alkali 
metals,  which  are  isomorphous  with  the  alums;  this  fact  illustrates 
the  close  resemblance  between  trivalent  manganese,  aluminium, 
chromium,  and  ferric  iron.  Potassium-manganic  sulphate  has  the 
formula  K2  SO4  ,  Mn2  (SO4)3  +  24  H2  O  (see  page  339). 

Manganous-manganic  oxide,  Mn3  04  ,  occurs  as  a  brownish-black, 
crystalline  mineral  known  as  hausmannite  ;  it  is  produced  by  heat- 
ing any  of  the  other  oxides  of  manganese,  when  in  contact  with 
the  air,  to  a  red  heat.*  Manganous-manganic  oxide  is  considered 
as  being  a  manganous  salt  of  or^Ao-manganous  acid  (the  hydroxide 
of  manganese  dioxide,  Mn  (OH  )4  ,  being  designated  as  manganous 
acid).-\  This  theory  is  expressed  by  the  following  formula:  — 


o 

n      =     n 

Mn 

Mn(OH)4-f  2MnO  =  Mn304     +  2H20; 


Mn      ;       +  2MnO  =  Mn^  +  2H20; 

[OH  [o[ 


and  is  borne  out  by  the  fact  that  manganoms-manganic  oxide,  when 
treated  with  dilute  nitric  or  sulphuric  acid,  forms  manganous  nitrate 
or  sulphate,  while  manganese  dioxide  is  left  behind  :  — 

1.  Mn  04  Mn2  +  4  HN03  =  Mn  (OH  )4  +  2  Mn  (  N03)2  . 

2.  Mn(OH)4  = 


This  reaction  is  similar  to  that  encountered  with  Pb3  04  (see  page 
323). 

Manganese  dioxide,  Mn  02  ,  is  probably  the  most  important  com- 
pound of  manganese.  It  occurs  in  large  quantities  as  a  mineral 
which  is  named  pyrolusite.  The  latter  has  a  steel-gray  color, 
metallic  lustre,  and  crystallizes  in  prisms  belonging  to  the  rhombic 
system.  A  considerable  amount  of  this  oxide  is  mined  in  the  New 
England  States,  and  in  California.  Manganese  dioxide  can  be  pre- 

*  The  oxide  Mn2  O3  is  stable  when  in  an  atmosphere  of  oxygen,  if  the 
temperature  is  no  higher  than  that  of  a  Bunsen  burner.  At  white  heat  it  is 
also  converted  into  Mn3  O4  . 

t  Manganous  acid  would  thus  be  parallel  with  sulphurous  acid  :  —  SO2  , 
sulphur  dioxide;  H2SO3,  sulphurous  acid;  H4SO4,  orthosulphurous  acid. 
Mn  O2  ,  manganese  dioxide  ;  H2  Mn  Os  ,  manganous  acid  ;  H4  Mn  O4  ,  ortho- 
manganous  acid. 


464  MANGANESE   DIOXIDE;    REACTIONS. 

pared  artificially  by  oxidizing   manganous    carbonate  with   an  al- 
kaline solution  of  chlorine  (see  pages  122  and  123). 

Manganese  dioxide,  when  heated  to  a  high  red  heat,  loses  one- 
third  of  its  oxygen,  and  changes  into  nianganous-manganic  oxide 
(see  page  19)  :  - 


During  this  decomposition,  the  dioxide  first  changes  into  manganic 
oxide,  Mn2  03  ,  and  then  the  latter  compound  loses  the  quantity 
of  oxygen  necessary  to  produce  Mn3  04  ,  as  the  temperature  is  in- 
creased to  a  high  red  heat.  Acids  decompose  manganese  dioxide. 
When  acting  in  the  cold  they  not  infrequently  produce  manganic 
salts,  while  the  surplus  of  oxygen  is  liberated  (see  page  462)  ;  on 
the  other  hand,  hot  acids  leave  manganous  salts  behind.  Of  course, 
if  any  oxidizable  substance  is  present,  the  liberated  oxygen  does  not 
pass  off  as  such,  but  is  used  up  in  the  work  of  oxidation.  The 
reaction,  when  warm  sulphuric  acid  is  brought  in  contact  with  man- 
ganese dioxide,  is  as  follows  :  — 

Mn  02  +  H2  S04  =  Mn  S04  +  H2  0  +  0, 

but,  on  the  other  hand,  hydrochloric  acid,  because  it  is  readily 
oxidized,  liberates  chlorine  :  — 

Mn  02  +  4  H  Cl  =  Mn  C12  +  2  H2  0  -f  2  Cl. 

In  the  latter  case  it  is  not  at  all  improbable  that  Mn  C14  is  at  first 
formed,  and  that  the  latter  salt  subsequently  breaks  down  into  man- 
ganous chloride  and  chlorine  (see  pages  60  and  61).*  The  reactions 
of  manganese  dioxide  are  very  similar  to  those  of  the  corresponding 
compound  of  lead  (see  page  323).  Several  hydroxides  related  to 

*  Recent  investigations  render  it  probable  that  the  chloride  Mn  C14  ,  when 
formed,  assumes  the  part  of  an  acidic  anhydride,  and,  uniting  with  the  excess 
of  hydrochloric  acid  which  is  present,  forms  an  acid  of  the  formula  H2  Mn  C16  , 
analogous  to  H0  Si  F6  ,  and  to  other  similar  compounds  which  we  have  encoun- 
tered (see  pages  303,  316  and  330).  The  compound  H2MnCl6  then  breaks 

down  as  follows  :  — 

H2  Mn  C16  =  2  H  Cl  +  Mn  C12  +  2  Cl. 

It  is,  however,  very  certain  that  the  compound  is  not  alone  present  in  the 
beginning  of  the  reaction  between  hydrochloric  acid  and  manganese  dioxide, 
for  manganese  trichloride,  Mn  C13  ,  is  also  produced  :  — 

MnO2  +  4  H  Cl  =  Mn  C13  +  Cl  +  2  H2  O. 
The  whole  matter  may  therefore  be  regarded  as  not  as  yet  definitely  settled. 


MANGANATES.  465 

manganese  dioxide  are  known.  The  simplest  of  these  is  Mn  0  (OH  )2. 
These  hydroxides  have  acidic  properties,*  and  form  salts  which  are 
designated  as  manganites.  However,  none  of  the  latter  are  derived 
from  the  simple  ortho-  or  meta-hydroxides  —  Mn  (OH  )4 ,  or 
Mn  0  (OH  )2  —  but,  like  the  salts  of  so  many  acids  which  we  have 
already  studied,  they  are  produced  by  complicated  poly-manganous 
acids.  Two  examples  of  the  latter  are  H2  Mn2  05  and  H2  Mn5  On  ; 
their  formation  might  be  imagined  as  taking  place  as  follows :  — 

1.  2MnO(OH)2  =  H2Mn205  +  H2  0. 

2.  5  Mn  0  (OH  )2  =  H2  Mn5  On  +  4  H2  0. 

POTASSIUM  PENTAMANGANITE,  K2  Mng  On ,  is  formed  by  passing  carbon 
dioxide  into  a  solution  of  potassium  manganate. 

CALCIUM  DIMANGANITE,  Ca  Mn2  O5 ,  is  important  because,  when  treated 
with  hydrochloric  acid,  it  generates  chlorine :  — 

Ca  O,  2  Mn  O2  +  10  H  Cl  =  Ca  C12  +  2  Mn  C12  +  5  H2  O  +  4  Cl.  This 
salt  is  readily  produced  by  heating  a  mixture  of  manganous  hydroxide 
and  calcium  hydroxide  in  a  current  of  air,  so  that  a  method  t  of  util- 
izing the  waste  manganous  chloride,  which  was  formerly  lost  during 
the  commercial  preparation  of  chlorine,  has  been  founded  on  this 
reaction. 

When  manganese  dioxide  is  fused  with  a  caustic  alkali  in  the 
presence  of  an  oxidizing  agent,  or  even  in  the  air,  a  manganate  is 
produced.  The  manganates,  in  chemical  composition,  are  analogous 
to  the  sulphates,  chromates,  and  molybdates,  etc.  This  will  be  seen 
from  the  following  formulae  :  — 

K2Mn04,  potassium  manganate. 
K2CrO4,  potassium  chromate. 
K2  Mo  04 ,  potassium  molybdate. 

However,  those  manganates  which  are  soluble  in  water  differ  very 
markedly  from  their  prototypes  in  one  particular;  they  are  very 
readily  decomposed  by  the  addition  of  an  excess  of  the  solvent  and 
are  stable  only  in  alkaline  solution.  In  so  decomposing,  they  change 
into  permanganates  and  manganese  dioxide.  The  manganates,  when 

*  The  hydroxide  of  manganese  dioxide,  MnO(OH)2,  is  sufficiently  acid 
in  its  properties  to  redden  blue  litmus  paper,  and  to  expel  carbonic  acid  from 
the  carbonates  of  the  alkalies.  Several  of  the  hydroxides  are  found  as  min- 
erals; they  are  termed  "  wad." 

t  Weldon's  process. 


466  PERMANGANIC   ACID. 

dry,  have  a  deep  red  color,  with  a  metallic  lustre.  They  are  power- 
ful oxidizers,  and  the  salts  of  the  alkali  metals  only  are  soluble  in 
water.  The  solutions  are  green,  but  change  to  red  on  addition  of 
an  acid.*  Neither  manganic  acid,  H2Mn04,  nor  its  anhydride, 
Mn  08 ,  has  been  isolated. 

The  anhydride  of  permanganic  acid,  Mn2  07 ,  is  the  only  one  of 
'the  oxides,  which  is  the  anhydride  of  an  acid  of  the  general  formula 
H  X  04  (and  which  consequently  is  related  to  perchloric  and  per-iodic 
acids),  which  has  been  isolated.  It  is  a  dark-green  (almost  black), 
oily  liquid,  produced  by  adding  potassium  permanganate,  in  small 
quantities,  to  concentrated  sulphuric  acid.  The  liquid  must  be 
cooled  by  means  of  a  mixture  of  snow  and  salt  during  the  process, 
and,  after  the  operation  is  completed,  warmed  to  ^60°,  when  the 
anhydride  distils.  The  highest  oxide  of  manganese  is  extremely 
unstable  ;  if  allowed  to  stand,  it  spontaneously  liberates  oxygen 
and  leaves  manganic  oxide,  Mn2  03 ;  it  is  a  most  powerful  oxidizer ; 
paper  and  alcohol  are  instantly  ignited  by  it.  When  added  to  water, 
it  forms  permanganic  acid  :  — 

Mn2  07  +  H2  0  =  2  H  Mn  04 . 

Permanganic  acid  is  entirely  analogous  to  perchloric  acid.  It 
can  be  produced  by  decomposing  a  solution  of  barium  permanganate 
with  exactly  the  requisite  amount  of  sulphuric  acid.  By  means  of 
the  ensuing  double  decomposition,  insoluble  barium  sulphate  and 
permanganic  acid  are  produced  :  — 

Ba  (Mn  04)2  +  H2  S04  =  Ba  S04  +  2  H  Mn  04; 

the  red  solution  so  formed  is  then  evaporated  to  dryness,  when  per- 
manganic acid  remains  in  the  form  of  a  reddish-brown,  crystalline 
substance.  Permanganic  acid  is  quite  unstable ;  it  breaks  down 
when  exposed  to  the  light,  and  its  solutions,  like  those  of  perchloric 
acid,  are  powerful  oxidizers. 

*  Due  to  the  formation  of  a  permanganate :  — 

5  K2  Mn  O4  +  4H2  SO4  =  4  K  Mn  O4  +  Mn  SO4  +  3  K2  SO4  +  4  H2  O. 
When  water,  and  not  acid,  is  added  to  the  manganate,  a  hydroxide  derived 
from  Mn  O2  is  formed :  — 

3  K2  Mn  O4  +  3  H2  O  =  2  K  MnO4  +  Mn  O  (OH)2  +  4  KOH. 
The  potassium  hydroxide  will  then  react  with  the  hydroxide  of  manganese  to 
form  a  manganite.     Very  weak  acids  (such  as  carbonic  acid)  facilitate  the 
change. 


PERMANGANATE    OF  POTASSIUM.  467 

The  most  important  permanganate  is  the  permanganate  of  potas- 
sium. This  salt  is  produced  by  fusing  manganese  dioxide  with  a 
mixture  of  potassium  hydroxide  and  an  oxidizing  salt  (such  as 
potassium  nitrate  or  potassium  chlorate  *)  ;  the  dark-green  flux 
then  contains  potassium  manganate,  which  is  converted  into  the 
permanganate  by  dissolving  in  water,  and  then  passing  carbon 
dioxide  into  the  solution,  f  Potassium  permanganate  crystallizes  in 
long  prisms,  belonging  to  the  monoclinic  system;  the  crystals  are 
dark  green,  almost  black,  in  color;  their  solution  in  water  has  an 
intense  reddish-purple  color  ;  the  salt  is  a  most  powerful  oxidizing 
agent.  In  oxidizing  with  potassium  permanganate,  there  is  an 
essential  difference  between  the  action  of  the  salt  in  acid  or  in 
alkaline  solution  ;  in  the  former  event  the  permanganate  changes 
to  a  manganous  salt;  in  the  latter  to  manganese  dioxide,  subse- 
quently to  a  manganite. 

1.     Acid  solution. 


2KMn04+3H2S04==K2S04  +  2MnS04  +  50+3  H20. 
2.     Alkaline  solution. 

a.  2  KMnO^rf  2  KOH  =  2  K2  Mn  0/+  H2  0  +  0  ; 

b.  2K2Mn04<f  2  H20=4  KOH  +  2  Mn02  +  2  0. 

The  first  change,  in  alkaline  solution,  is,  therefore,  from  the  per- 
manganate to  the  manganate,  the  second  from  the  manganate  to 
manganese  dioxide;  of  course,  the  latter  substance  subsequently 
produces  a  manganite  with  the  excess  of  caustic  potash  which  is 
present.  From  these  equations  it  follows  that,  for  every  two  for- 
mula-weights of  potassium  permanganate  in  acid  solution,  there 
are  five  atoms  of  oxygen  to  be  used  in  oxidation,  while  for  every 
two  formula-weights  in  alkaline  solution,  there  are  but  three.  The 
following  two  equations  illustrate  the  application  of  these  rules  :  — 

2  K  Mn  04  +  3  H2  S04  +  5  H2  S03  =  K2  S04  +  2  Mn  S04  + 

5  ff2S04  +  3  H20. 
2  KMn04  +  8  Na2S03  +  H20  =  2  KOH  +  2  Mn02  + 


*  The  oxygen  of  the  atmosphere  is  also  able  to  effect  the  change. 

t  The  reaction  is  as  follows  :  — 

3  K2  Mn  O4  +  2  CO2  =  2  K2  CO3  +  2  K  Mn  O4  +  Mn  O2; 

the  reactions  taking  place  on  the  addition  of  water  or  of  sulphuric  acid  are 
given  on  page  466.  foot-note. 


468  MANGANESE;    COMPOUNDS   OF. 

Owing  to  its  oxidizing  powers,  potassium  permanganate  is  frequently 
used  as  a  disinfecting  agent. 

Manganese,  with  its  diversity  of  compounds  in  which  it  displays 
such  a  great  difference  in  valence,  serves  admirably  to  illustrate  the 
fact  that  the  chemical  behavior  of  an  element  is,  to  a  large  extent, 
relative,  and  that  the  properties  of  its  compounds  depend  just  as 
much  upon  the  elements  with  which  it  is  united,  and  upon  the  manner 
of  such  union,  as  they  do  upon  the  individual  characteristics  of  the 
element  itself.  Manganese,  when  entering  into  the  formation  of 
manganous  salts,  is  far  more  like  magnesium  or  zinc  than  it  is  like 
manganese  in  the  inanganates  or  permanganates  ;  when  manganic 
oxide,  as  a  base,  has  united  with  acids  to  form  manganic  salts,  then 
manganese  resembles  trivalent  chromium,  iron,  or  aluminium,  even 
to  such  an  extent  that  its  sulphate  will  form  alums  ;  in  the  man- 
ganates  the  element  may  be  compared  to  sulphur  in  the  sulphates  ; 
while,  lastly,  the  permanganates  are  analogous  to  the  perchlorates 
or  per-iodates.  One  essential  distinction,  however,  remains  as  exist- 
ing between  the  various  compounds  of  manganese  and  the  elements 
with  which  it  has  been  compared  :  the  compounds  of  manganese  can 
be  converted  the  one  into  the  other,  while,  of  course,  those  of  zinc, 
for  example,  can  never  be  changed  into  those  of  chromium,  or  those 
of  iron  into  those  of  chlorine.  Oxidizing  agents  transform  manga- 
nous salts  into  manganic  salts,  manganese  dioxide,  manganates,  and 
permanganates  successively  ;  while,  on  the  other  hand,  reducing 
agents,  beginning  with  the  permanganates,  produce  exactly  the 
opposite  result. 

The  relationship  between  the  compounds  of  manganese  and  those 
of  a  few  other  elements  may  be  seen  from  the  following  table  :  — 

Manganous  compounds  (  Sulphate,  Mn  SO4  *  7  H2  O  )  (  Zn  SO4  +  7  H,  O 

(Oxide  Mn<m          \  Nitrate,  Mn  (NO3)2  >  l*e  <  Mg(NO,)a 

(  Chloride,  Mn  C12  )          (  CaCl2 

Manganic  compounds  (  Sulphate,  Mna  (SO4)3 

(Oxide,  Mn2  O,)       )  Alum,  Mn3  (SO4)3  K2  SO4  ,  24  H2  O 

{. 

Manganates        (  Potassium  manganate,  K2Mn  O4  |  like  i  K2  SO4  or  K2  Cr  O4 
(Acid,  H2MnO4)  (  Barium  manganate,  Ba  Mn  O4       f          \  BaSO4  or  BaCrO4 

Permanganates  (  Potassium  permanganate,  K  Mn  O4    |  j..      (  K  Cl  O4 
(Acid,  HMnO4)  (  Barium  permanganate,  Ba(MnO4)2  j  (  Ba(ClO4)2 

The  oxides  of  manganese  are,  perhaps,  most  like  those  of  lead  ; 
but,  in  formula,  they  also  resemble  those  of  the  type  of  the  family, 
chlorine. 


... 
llke 


MANGANESE;    COMPOUNDS   OF.  469 

OXIDES   OF    CHLORINE.  LEAD.  MANGANESE. 

C120 

Cl  0                                  Pb  O  Mn  O 

C12  0  3                               Pb2  08  Mn2  03 

C102                                  Pb02  Mn02 

(C1207)  'Mn207 

Pb304  Mn304 

This  process  of  comparison,  were  space  to  permit,  could  be  car- 
ried much  farther ;  and,  indeed,  the  formation  of  tables  like  the 
above  would  be  a  most  instructive  exercise  for  the  pupil. 


470  IRON  ;   COBALT  ;   NICKEL. 


CHAPTER   LIX. 

IRON,    COBALT,   AND   NICKEL. 

Iron  ;  symbol,  Fe  ;  atomic  weight,  56  ; 
Cobalt ;  symbol,  Co ;  atomic  weight,  59  ; 
Nickel ;  symbol,  Ni ;  atomic  weight,  58.7. 

IRON,  cobalt,  and  nickel  are  members  of  the  eighth  family  of 
elements,  which  consists  of  three  groups,  each  of  which  contains 
three  individuals,  namely :  — 

1.  Iron,  cobalt,  nickel. 

2.  Euthenium,  rhodium,  palladium. 

3.  Osmium,  iridium,  platinum. 

The  properties  of  these  elements  are  such  that  they  form  a 
gradual  transition  from  the  last  elements  of  the  first  halves  of  the 
long  periods,  to  the  first  ones  in  the  second  halves  ;  so  that  we  should 
expect  iron  to  be  very  much  like  manganese,  and  nickel  to  bear 
marked  resemblance  to  copper ;  and  such  is  the  case.  The  valence 
of  the  elements  in  their  highest  oxides,  passing  from  manganese 
(through  iron,  cobalt,  and  nickel)  to  copper,  diminishes  with  each 
individual  as  the  atomic  weight  increases  ;  manganese,  in  the  per- 
manganates, has  a  valence  of  seven ;  iron,  in  ferric  acid,  a  valence 
of  six ;  cobalt  a  valence  of  three  in  the  oxide  Co2  03 ;  *  nickel,  almost 
without  exception,  forms  compounds  derived  from  Ni  0 ;  while, 
lastly,  copper  can  appear  as  a  univalent  metal  in  its  cuprous  form. 

Mn.  Fe.  Co.  Ni.  Cu. 

Highest  valence,  VII  VI  III          II  t        I  ( II ) 

Oxides,  Mn207      Fe2  06        Co2  03      Ni2  02      Cu2  0. 

Iron,  cobalt,  and  nickel  are  near  the  minimum  of  the  curve  of 

*  An  oxide  of  cobalt,  Co  O2 ,  probably  also  exists. 

t  Nickel  can  form  an  oxide  Ni2  O3 ;  but  the  latter  forms  no  salts,  and  is 
decomposed  with  the  greatest  of  ease. 


IRON  ;    COBALT  ;    NICKEL. 


471 


atomic  volumes  formed  in  the  period  of  which  they  are  members, 
while  the  elements  which  follow  in  the  same  period  show  a  rapid 
increase  in  their  atomic  volumes  as  we  pass  along  the  series  in  the 
direction  of  increasing  atomic  weights ;  the  three  individuals  in 
question  are  therefore  malleable  and  ductile,  have  high  melting 
points,*  and  form  colored  salts.  The  physical  constants  mentioned 
in  this  connection  are  given  in  the  following  table :  — 


ATOMIC  WEIGHT. 

SPECIFIC  GEAVITY. 

ATOMIC  VOLUME. 

MELTING  POINT. 

Iron 

56. 

7.8 

7.2 

1770°  (?) 

Cobalt 

59. 

8.5 

6.9 

1750° 

Nickel 

58.7 

8.8 

6.7 

1570° 

Nickel,  owing  to  its  chemical  reactions,  specific  gravity,  atomic 
volume,  and  melting  point,  apparently  has  its  position  in  the  periodic 
system  immediately  following  that  of  cobalt,  although  its  atomic 
weight  is  somewhat  less  than  that  of  the  latter  element ;  it  seems 
probable,  therefore,  that,  at  some  future  time,  more  exact  study  will 
prove  the  atomic  weights  of  the  two  elements  in  question  to  have 
been  inaccurately  determined ;  f  if  this  should  not  prove  to  be  the 
case,  however,  then  cobalt  and  nickel  certainly  form  a  most  remark- 
able exception  to  Meiidelejeff's  rule,  in  the  arrangement  of  that 
system. 

The  principal  minerals  in  which  iron,  cobalt,  and  nickel  occur 
are  as  follows  :  — 

*  The  melting  points  are,  however,  lower  than  those  of  the  elements  im- 
mediately preceding  which  have  diminishing  atomic  volumes  with  increasing 
atomic  weights. 

t  Gerhardt  Kriiss  has  recently  published  some  work  in  which  he  under- 
takes to  show  that  what  has  hitherto  been  regarded  as  pure  nickel  in  reality 
contains  an  admixture  of  one,  or  more,  hitherto  undiscovered  elements,  and 
that  the  same  is  probably  true  of  cobalt.  Some  of  the  fractions  into  which 
Kriiss  divided  nickel  have  an  atomic  weight  lying  between  56  and  58,  the  others 
between  60  and  100;  in  view  of  these  results,  the  atomic  weights  of  cobalt  and 
nickel  are  as  yet  undetermined.  Clemens  Winckler,  however,  by  reason  of  his 
previous  investigations  on  the  atomic  weight  of  nickel  and  by  reason  of  a 
review  on  some  parts  of  his  former  work,  which  he  instituted  with  apparently 
pure  materials,  doubts  Kriiss's  results;  so  that,  until  further  light  is  thrown  on 
the  subject,  the  old  theories  as  regards  cobalt  and  nickel  must  be  maintained. 
See  Kriiss  and  Schmidt ;  Berichte  d.  Deutsch.  Chem.  Gesell. ;  22,  11  and  2026  ; 
Clemens  Winckler,  ibid.  890.  See  also  Mond,  Journ.  Chem.  Soc.  1890,  753. 
Mond  has  prepared  chemically  pure  nickel. 


472  IRON  ;   COBALT  ;   NICKEL  ;   OCCURRENCE. 

Native  iron.  The  occurrence  of  masses  of  iron  of  terrestrial  origin  has 
been  mentioned  several  times,  but  is  not  beyond  doubt.  Meteoric  iron 
is  not  infrequently  found;  it  usually  contains  from  1  to  20  per  cent  of 
nickel.  These  meteorites  contain  the  metal  arranged  in  striae  with  a  dif- 
fering contents  of  nickel,  so  that,  as  they  offer  a  differing  resistance  to 
acids,  meteoric  iron,  when  polished  and  subjected  to  the  corroding 
action  of  reagents,  will  show  a  surface  marked  by  regular  etchings. 

Iron  pyrites,  Fe  82 ,  occurs  in  rocks  of  all  ages :  it  is  isomorphous  with  arsen- 
ical pyrites,  FeAsS,  with  the  sulphide  of  manganese,  MnS2  (haiie- 
rite),  and  with  the  sulphides  and  arsenides  of  cobalt  and  nickel,  having 
the  general  formulae  MS2 ,  M  As2 ,  or  M  As  S.  Iron  pyrites  is  dimor- 
phous, for  a  mineral  of  the  same  formula,  belonging  to  a  different  crys- 
talline system,  is  known;  this  mineral  is  called  markasite. 

Ferrous  sulphide,  Fe  S,  is  found  as  troilite. 

Ferric  sulphide,  Fe2  83  ,  frequently  plays  the  part  of  an  acidic  anhydride, 
and  with  bases,  such  as  Cu2  S,  Ag2  S,  or  Cu  S,  forms  minerals  of  which 
chalcopyrite,  Cu2  S,  Fe2  S3  =  2  Cu  Fe  S2 ,  is  an  example  (see  page  397). 

Hematite  (specular  iron)  is  ferric  oxide,  Fe2  Og .  It  is  one  of  the  most  im- 
portant iron  ores,  and  occurs  in  rocks  of  all  ages. 

Magnetite  (magnetic  iron  ore)  is  ferrous-ferric  oxide,  Fe  O,  Fe2  O3  =  Fe3  O4 . 
This  oxide  is  isomorphous  with  the  spinells  (page  339),  and  is  probably 
similarly  constituted. 

Various  hydroxides  of  ferric  oxide  are  also  frequently  met  with. 
The  chief  representative  of  this  most  important  class  of  minerals  is 
limonite  (  brown  hematite),  Fe4  09  H6  =  2  Fe2  03  -f  6  H2  0. 

Siderite  (spathic  iron)  is  ferrous  carbonate,  Fe  COs  .  It  occurs  in  many 
rock  strata,  in  gneiss,  mica  slate,  clay  slate,  and  with  the  coal  forma- 
tion. It  is  isomorphous  with  calcite. 

Iron  is  also  found  as  a  constituent  of  a  large  number  of  silicates, 
In  consequence  of  the  disintegration  of  the  rocks  in  which  it  occurs, 
it  finds  its  way  into  the  soil  and  into  the  natural  waters.  It  is  an 
invariable  constituent  of  chlorophyll  (the  green  coloring  matter  of 
leaves),  and  it  is  always  found  in  the  haemoglobin  of  the  blood. 

Cobalt  chiefly  occurs  as  cobaltite,  Co  As  S  (CoS.2  4-  Co  As2  ),  isomorphous 
with  iron  pyrites;  *  as  smaltite,  Co  (Fe  Ni)  As  2 ,  also  isomorphous  with 
iron  pyrites;  and  as  danaite,  (Fe,  Co)  (AsS)2,  isomorphous  with 
markasite.  Cobaltous  carbonate,  (sphserocobaltite)  CoCO3,isalso 

*  As  might  be  expected,  cobalt  replaces  iron  isomorphously;  but  in  this 
mineral  we  have  another  phenomenon  which  is  not  so  self-evident,  namely,  ar- 
senic replaces  sulphur  isomorphously.  This  substitution  is  not  infrequent  in 
the  group  of  minerals  of  which  iron  pyrites  is  the  representative  (see  page  234). 


.    . 

IRON  ;   METALLURGY.  473 

sometimes  found,  as  well  as  the  arsenate,  Co3  ( As  O4)2  +  8  H2  O,  which 
is  called  erythrite. 

Nickel  occurs  in  meteorites  as  an  alloy  of  iron ;  as  millerite,  Ni  S,  isomor- 
phous  with  niccolite,  as  gersdorfite,  Ni  As  S,  isomorphous  with  cobalt- 
ite  and  iron  pyrites;  as  the  arsenide,  niccolite,  Ni  As  (isomorphous  with 
zinc-blende,  page  426);  as  a  basic  carbonate,  as  the  arsenate,  and  as  a 
double  sulphide  of  iron  and  nickel  (Fe,  Ni)  S.  This  sulphide  occurs 
as  a  massive  variety,  and  is  termed  pentlandite. 

Iron,  cobalt,  and  nickel  are  all  easily  reduced  from  their  oxides 
by  means  of  charcoal.  The  metallurgy  of  iron  is  among  the  most 
important  commercial  operations  of  the  present  time.  In  the  prep- 
aration of  this  metal,  the  most  important  ores  are  the  oxides,  hy- 
droxides, and  the  carbonate.  The  ores  are  crushed  when  necessary, 
and  sometimes  roasted  for  the  purpose  of  expelling  the  water  and 
carbon  dioxide,  and  of  changing  the  oxides  as  much  as  possible  into 
ferric  oxide,  Fe2  03  .*  They  are  then  reduced  in  a  blast  furnace. 
The  latter  consists  of  a  shaft,  varying  in  height  from  fifty  to  ninety 
feet,  with  a  maximum  diameter  of  from  twenty  to  twenty-three  feet, 
the  shaft  being  shaped  like  two  truncated  cones  united  at  their 
bases ;  below  these  is  a  circular  chamber  or  hearth  which  is  built  of 
firebrick,  and  in  which  two  small  openings  can  be  made,  the  one 
higher  up  than  the  other  ;  the  lower  one  is  used  for  drawing  off  the 
molten  metal,  the  upper  one  for  the  slag.  The  remainder  of  the 
furnace  is  constructed  of  firebrick  and  encased  in  boiler  iron.  The 
blast  is  introduced  through  from  six  to  eight  openings,  termed  tuy- 
eres, at  from  four  to  six  feet  above  the  bottom  of  the  furnace,  and 
the  air  which  is  forced  in  through  these  is  heated  by  means  of  the 
waste  gases  passing  from  the  furnace.  The  furnace  is  charged  with 
alternate  layers  of  iron  ore,  coke  or  charcoal,  and  limestone.  The 
latter  substance,  uniting  with  the  siliceous  matter  f  which  is  present 
in  the  ore,  forms  a  fusible  glass  called  the  slag.$  The  molten  slag 
collects  at  the  bottom  of  the  furnace,  with  the  metal,  and,  being  of 

*  Some  sulphur-bearing  ores  are  roasted  to  burn  off  the  sulphur,  but  these 
ores  form  a  very  small  proportion  of  the  total  iron  compounds  used.  I  am 
indebted  to  Prof.  E.  D.  Campbell  for  a  review  of  the  pages  relating  to  the 
metallurgy  of  iron. 

t  Feldspar,  slate,  quartz,  etc. 

J  If  the  ore  contains  limestone  in  sufficient  or  excessive  quantity,  it  may 
be  necessary  to  add  siliceous  matter  such  as  feldspar.  The  lime  prevents  the 
formation  of  a  ferruginous  slag,  which  would  entail  a  loss  of  iron. 


474  IKON;    METALLURGY. 

less  specific  gravity,  floats  upon  the  surface  ;  it  can  therefore  be 
drawn  off  at  the  upper  small  opening  (termed  the  cinder  notch)  in 
the  base  of  the  furnace.  The  slag  is  not  formed  until  almost  com- 
plete reduction  has  taken  place,  so  that  it  exercises  its  protective 
action  only  upon  the  molten  metal  at  the  bottom  of  the  furnace. 

The  chemical  changes  which  take  place  in  a  blast  furnace  are 
quite  complicated,  and  all  of  them  are  not,  as  yet,  definitely  under- 
stood ;  however,  the  most  important  reactions  in  the  production  of 
cast  iron  are  as  follows.  The  carbon,  uniting  with  the  oxygen  en- 
tering from  the  tuyeres,  forms  carbon  dioxide  immediately  above 
those  openings,  but  the  carbon  dioxide  is  almost  instantly  and  com- 
pletely reduced  to  carbon  monoxide  in  passing  over  the  red-hot  coke 
or  charcoal  (see  page  287)  ;  hot  carbon  monoxide  now  comes  in  con- 
tact with  the  descending  charges  of  ore,  and  reduces  the  oxide  to  a 
spongy  form  of  iron  :  — 

Ee2  03  +  3  CO  =  2  Fe  +  3  C02 . 

The  portion  of  the  furnace  in  which  this  reduction  occurs  has  a  tem- 
perature of  from  450°  to  900°.  The  spongy  metal  passes  downward 
in  the  furnace,  the  temperature  increasing  to  950°  at  the  widest 
part  of  the  furnace ;  at  this  point  the  iron  takes  up  carbon  to  form 
a  chemical  combination  with  that  element,  and  at  a  lower  zone, 
when  the  temperature  is  about  1400°,  the  mass,  which  has  been  in 
a  pasty  condition,  melts  and  runs  down  into  the  hearth.  The  latter 
is  tapped  from  time  to  time,  and  the  iron  cast  into  semi-cylindrical 
moulds  called  "  pigs." 

Pig  iron  is  quite  impure ;  it  contains  carbon,  silicon,  sulphur, 
phosphorus,  and  manganese,  and  is  divided  into  two  chief  classes, 
white  cast  iron  *  and  gray  cast  iron  ;  the  former  contains  the  greater 
3>art  of  its  carbon  chemically  united  with  the  iron-,  the  latter,  in 
consequence  of  the  presence  of  silicon,  has  separated  the  major  por- 
tion of  its  carbon  in  the  form  of  graphite.  The  proportion  of  carbon 
in  white  iron  is  from  3  to  4.4  per  cent ;  in  gray  iron,  the  total 
carbon  is  about  the  same  as  in  white  iron,  but,  because  by  far  the 
greater  proportion  of  it  is  in  the  graphitic  for,m,  the  iron  is  of  a 
darker  color.  If  the  iron  ore  contained  a  considerable  quantity  of 
manganese,  the  latter  is  reduced  with  the  iron,  and  the  alloy  so 

*  White  cast  iron  is  formed  at  a  lower  furnace  temperature  than  gray  cast 
iron  or  spiegeliron. 


. 

WKOUG^IT  IRON  ;   STEEL. 

formed  is  capable  of  taking  up  a  considerably  greater  quantity  of 
carbon  (as  high  as  6.9  per  cent);  this  form  of  iron  is  known  as 
spiegeliron.  Cast  iron  is  brittle,  easily  fusible  (its  melting  point  is 
about  1050°  )  ;  it  cannot  be  welded  or  tempered  ;  when  treated  with 
hydrochloric  acid  it  dissolves,  and  the  combined  carbon  passes  off  in 
the  form  of  hydrocarbons,  which  possess  a  most  disagreeable  odor, 
while  the  graphite  remains  undissolved. 

Wrought  iron  is  produced  from  cast  iron  by  puddling.  Pud- 
dling consists  of  melting  cast  iron  in  a  furnace  lined  with  ferric 
oxide.  The  oxygen  of  this  lining,  combining  with  the  carbon,  sili- 
con, phosphorus,  and  manganese,  forms  the  oxides  of  those  elements. 
The  carbon  passes  off  as  carbon  monoxide ;  the  silicon  dioxide, 
phosphoric  acid,  and  manganous  oxide  unite  with  the  excess  of  the 
oxides  of  iron  to  form  a  slag,  so  that  a  metal  which  is  nearly  pure 
is  formed.  Wrought  iron  contains  less  than  .3  per  cent  of  carbon* 
and  very  small  amounts  of  silicon  and  phosphorus ;  it  possesses  a 
fibrous  texture,  is  malleable  and  ductile,  and  melts  at  about  1900°. 

Steel  contains  more  carbon  than  wrought  iron,  and  always  less 
than  cast  iron ;  it  is  produced  by  the  Bessemer,  open-hearth,  and 
crucible  processes. 

The  Bessemer  steel  process  produces  steel  directly  from  cast 
iron;  it  consists,  briefly,  of  first  burning  out  the  impurities  in 
melted  cast  iron,  by  placing  the  latter  in  a  large  converter  and 
forcing  air  in  through  the  bottom,f  and  then,  after  stopping  the 
blast,  of  adding  spiegeliron  $  until  the  requisite  amount  of  carbon  is 
present.  Bessemer  steel  is  used  in  the  manufacture  of  rails  and  of 
other  large  steel  implements ;  it  contains  from  .1  to  1.  per  cent  of 
combined  carbon. 

The  open-hearth  process  consists  in  melting  (in  large  regenerative 
furnaces)  a  mixture  of  steel  scraps  and  a  small  proportion  of  cast 
iron,  and  of  then  adding  a  sufficient  quantity  of  spiegeliron  to  the 
molten  metal  to  give  the  desired  percentage  of  carbon  and  man- 

*  When  it  contains  more  than  .3  per  cent  it  is  steel. 

t  The  oxidizing  action  of  the  air  maintains  the  high  temperature. 

%  If  the  cast  iron  contains  a  large  amount  of  phosphorus  the  crucibles 
(converters)  are  lined  with  burned  dolomite  or  magnesite  (see  page  416),  the 
phosphorus  is  then  oxidized,  and  forms  calcium  phosphate  with  lime  which  is 
added  with  the  metal.  This  latter  process  is  termed  the  "basic  Bessemer 
process." 


476  IRON;    PROPERTIES. 

ganese.     Open-hearth  steel  has  the  same  composition  as  Bessemer 
steel. 

Crucible  steel  is  made  by  melting  (in  small  covered  crucibles) 
purest  wrought  iron,  after  adding  a  sufficient  supply  of  charcoal, 
with  pure  pig  iron,  and  a  small  amount  of  manganese  dioxide. 
This  steel  contains  from  .6  to  2.  per  cent  of  carbon ;  and  the  purity 
of  the  materials  from  which  it  is  made,  as  well  as  its  freedom  from 
dissolved  oxides  and  gases,  renders  this  class  of  steel  the  only  one 
suitable  for  the  manufacture  of  the  highest  grade  of  tools. 

When  steel  is  heated,  and  rapidly  cooled  by  plunging  into  water, 
it  becomes  very  hard  and  brittle ;  this  hardened  steel,  when  once 
more  heated  and  allowed  to  cool  slowly,  becomes  elastic ;  this  pro- 
cess is  called  tempering.  Steel  is  capable  of  taking  a  very  high 
polish. 

Chemically  pure  iron  can  be  prepared  by  reducing  either  the  pure 
oxide,  oxalate,  or  chloride  of  iron  in  a  current  of  hydrogen ;  it  has  a 
specific  gravity  of  7.84,  and  a  melting  point  which  is  probably  not 
less  than  1800° ;  *  it  is  bluish-gray,  almost  white,  in  color,  and  is 
malleable  and  ductile ;  one  of  the  most  striking  physical  properties 
that  it  possesses  is  that  of  magnetism.  Pure  iron  loses  its  magnet- 
ism as  soon  as  a  magnet,  which  has  been  placed  in  its  neighbor- 
hood, is  removed  ;  steel,  however,  is  able  to  retain  the  property. 
Pure  iron  is  not  attacked  by  dry  oxygen  at  ordinary  temperatures ; 
when  exposed  to  moist  air  it  undergoes  slow  oxidation, f  forming 
ferric  oxide,  Fe2  03 ;  if  it  is  heated  in  oxygen  it  burns,  forming 
Fe3  04 ,  mixed  with  Fe2  03  (see  page  22) ;  it  unites  with  the  halo- 
gens in  the  same  manner.  Iron  will  rust  when  placed  under  water 
which  contains  dissolved  oxygen ;  this  action  is  accelerated  by  the 
presence  of  acids  and  retarded  by  .the  presence  of  alkalies.  $  Dilute 
hydrochloric  or  sulphuric  acids  dissolve  iron,  liberating  hydrogen 
and  forming  ferrous  chloride  and  ferrous  sulphate  respectively  (see 
page  32)  ;  concentrated  sulphuric  acid,  when  cold,  is  without  action, 

*  The  melting  point  of  pure  iron  has  been  variously  given  as  being  1560°, 
1587°,  1600°,  1800°,  and  it  has  even  been  stated  that  absolutely  pure  iron  is 
infusible. 

t  It  is  stated  that  moist,  pure  air  does  not  attack  iron ;  the  rusting  tak- 
ing place  only  if  carbon  dioxide  is  present. 

J  The  rusting  is  also  assisted  by  the  presence  of  salts,  especially  of  those 
of  ammonium. 


COBALT;  NICKEL;  METALLURGY.  477 

when  heated  with  iron  it  liberates  sulphur  dioxide  and  produces 
ferric  sulphate  (see  page  136)  ;  concentrated  nitric  acid  has  the  same 
effect  on  iron  as  it  has  on  aluminium  (see  page  334),  the  metal  does 
not  dissolve  but  is  transferred  to  the  "  passive  state ;  "  when  in  this 
condition  it  is  no  longer  attacked  by  the  dilute  nitric  acid,  nor  will 
it  separate  copper  from  a  solution  of  copper  sulphate,  a  reaction 
into  which  ordinary  iron  very  readily  enters  :  — 

Cu  S04  +  Fe  =  Fe  S04  +  Cu.* 

Several  explanations  as  to  the  reason  of  this  condition  have  been  of- 
fered. One  of  these  is  that  the  iron  becomes  covered  with  a  very 
thin  layer  of  ferrous-ferric  oxide,  which  is  insoluble  in  nitric  acid  ; 
this  theory,  however,  is  without  absolute  experimental  proof,  f 
Dilute  nitric  acid  dissolves  iron,  forming  ferrous  nitrate,  while  a  por- 
tion of  the  acid  is  reduced  to  ammonium  nitrate  (see  page  206,  a). 

Cobalt  is  quite  difficult  to  obtain  in  a  pure  state,  because  its  ores  always 
contain  iron  and  nickel,  from  which  latter  element  the  metal  is  not  easy 
to  separate;  copper,  bismuth,  lead,  or  silver  may  also  be  present.  The 
chief  points  in  the  separation  are :  first,  the  burning  away  of  the  sul- 
phur and  arsenic  present  in  the  ores;  and  secondly,  the  separation  of 
the  copper,  bismuth,  etc.,  by  means  of  sulphuretted  hydrogen,  after 
the  oxides  produced  by  the  roasting  have  been  dissolved  in  acids. 
The  solutions  which  remain  after  the  sulphides  (which  have  been 
precipitated)  have  been  filtered  off,  are  treated  with  chlorine  and  cal- 
cium hypochlorite ;  by  this  means  the  cobalt  salts,  which  are  present, 
are  oxidized  to  insoluble  cobaltic  oxide  before  those  of  nickel  are 
attacked;  the  cobaltic  oxide,  when  the  operation  has  gone  just  far 
enough,  is  separated,  dried,  and  reduced  to  metallic  cobalt  by  heating 
in  a  current  of  hydrogen. 

Nickel  is  more  easily  obtained  than  cobalt,  because  it  is  present  in  greater 
quantity,  and  because  its  ores  are,  as  a  rule,  purer.  The  sulphides  or 
arsenides  of  nickel  are  roasted  in  the  same  manner  as  those  of  cobalt ; 
the  oxide  so  obtained  is  mixed  with  charcoal  and  heated,  by  which 
means  reduction  to  metallic  nickel  takes  place. 

Cobalt  forms  crystalline  metallic  plates  which  have  a  specific 
gravity  of  8.5,  and  which  melt  at  a  somewhat  lower  temperature 

*  See  page  313,  foot-note. 

t  In  support  of  this  theory  is  the  fact  that  the  passive  state  can  also  be 
produced  by  other  oxidizing  agents,  such  as  chloric,  bromic,  iodic,  or  chromic 
acids.  Another  theory  of  considerable  plausibility  is  that  the  iron  becomes 
covered  with  a  thin  layer  of  gas;  both  hypotheses  are  borne  out  by  the  fact 
that  passive  iron,  when  rubbed,  returns  to  its  normal  state. 


478  COBALT;    NICKEL;    PROPERTIES. 

than  iron.  Like  the  latter,  cobalt  is  capable  of  attracting  the 
magnet.  The  metal  is  susceptible  of  a  very  high  polish,  is  malle- 
able and  very  ductile.  The  metal,  after  it  has  been  cast  into  solid 
pieces,  is  entirely  unaltered  by  exposure  to  the  air ;  at  white  heat, 
however,  it  burns  to  form  cobaltous-cobaltic  oxide,  Co3  04 .  When 
heated  and  then  plunged  into  concentrated  nitric  acid  it  becomes 
"passive."  Hydrochloric  or  sulphuric  acids  slowly  dissolve  the 
metal,  forming  cobaltous  chloride  and  cobaltous  sulphate  respectively. 
Dilute  nitric  acid  readily  dissolves  cobalt  to  produce  cobaltous 
nitrate. 

Nickel.  It  is  somewhat  doubtful  if  chemically  pure  nickel  has 
ever  been  obtained*  (see  page  471)  ;  that  which  has  hitherto  been 
regarded  as  such  is  produced  by  reduction  from  the  pure  oxalate 
of  nickel.  When  cast  into  cubes  it  is  a  lustrous,  almost  silver- 
white  metal  which  is  nearly  as  malleable  and  ductile  as  iron.  It 
melts  at  a  temperature  lower  than  the  melting  point  of  either 
cobalt  or  iron.  It  is  attracted  by  the  magnet  at  ordinary  temper- 
atures, but  loses  this  property  when  heated  to  350°.  Nickel  which 
has  been  cast  into  solid  pieces  is  not  oxidized  in  the  air,  and  it 
scarcely  burns  even  when  heated  white  hot  in  an  atmosphere  of 
oxygen.  The  metal  is  but  slowly  attacked  by  hydrochloric  or  sul- 
phuric acid;  it  readily  dissolves  in  nitric  acid  to  form  nickelous 
nitrate.  Concentrated  nitric  acid  renders  the  metal  "passive." 
The  specific  gravity  of  nickel  is  8.9. 

Alloys  of  iron.  Alloys  of  manganese,  tungsten,  or  nickel,  with 
steel,  possess  great  toughness,  and  are  of  growing  commercial  im- 
portance. The  power  which  zinc  and  tin  possess,  of  adhering  to  the 
surfaces  of  iron  sheets,  is  probably  due  to  the  formation  of  alloys. 

Alloys  of  nickel  are  used  in  the  preparation  of  coins,f  in  the 
manufacture  of  German  silver  (which  contains  copper,  zinc,  and 
nickel)  and  in  a  number  of  other  ways.  Nickel-plating  is  accom- 
plished by  electrolyzing  a  solution  of  nickel-ammonium  sulphate,  | 
the  metal  to  be  covered  being  the  negative  electrode. 

Compounds  of  iron.     Iron  forms  two  series  of  compounds,  fer- 

*  Mond,  Journ.  Chem.  Soc.  1890,  753,  asserts  that  chemically  pure  nickel 
can  be  prepared  by  the  decomposition  of  a  liquid  compound  of  nickel  and  carbon- 
monoxide,  Ni(CO)4.  The  atomic  weight  of  this  nickel  is  58.69. 

t  The  nickel  coins  of  the  United  States  contain  75%  copper  and  25%  nickel. 

J  Steel  can  be  nickel-plated  by  simply  plunging  into  a  bath  of  nickel-ammo- 
nium sulphate. 


FERROUS   COMPOUNDS.  479 

rous  compounds,  in  which  the  metal  is  divalent,  and  ferric  com- 
pounds, in  which  it  is  trivalent.  In  addition  to  these  two  stages  of 
oxidization,  there  exists  a  ferric  acid  (the  anhydride  of  which  would 
be  Fe  08),  in  which  iron  is  hexavalent. 

Ferrous  oxide,  Fe  0,  is  produced  by  reducing  ferric  oxide  in  a 
current  of  hydrogen ;  the  corresponding  hydroxide  has  the  formula 
Fe  (OH)2.  This  compound  is  precipitated  by  adding  ammonia 
water  to  a  solution  of  a  ferrous  salt ;  the  precipitate  is  white,  but 
it  rapidly  turns  brown  on  exposure  to  the  air,  because  it  changes 
into  ferric  hydroxide.  Both  ferrous  oxide  and  hydroxide  are  bases ; 
they  dissolve  in  acids  to  form  ferrous  salts. 

FERROUS  CHLORIDE  is  produced  by  dissolving  iron  in  hydrochloric  acid, 
while  excluding  the  air.  The  dry  salt,  having  the  formula  Fe  C12 ,  can 
then  be  isolated  by  evaporating  the  solution  in  a  current  of  hydrogen. 
It  is  a  white,  crystalline  mass  which  volatilizes  at  a  high  red  heat.  The 
vapor  density  of  ferrous  chloride  at  white  heat,  air  =  1,  is  4.39.  This 
number  corresponds  to  a  molecular  weight  given  by  the  formula 
Fe  G12 .  At  a  lower  temperature  the  molecules  are  probably  Fe2  C14  . 
When  exposed  to  the  air,  ferrous  chloride  rapidly  oxidizes  to  a  mix- 
ture of  ferric  chloride  and  ferric  oxide :  — 

6  Fe  C12  +  3  O  =  4  Fe  C13  +  Fe2O3 . 

When  in  solution,  provided  hydrochloric  acid  is  present,  ferric  chlo- 
ride alone  is  produced :  — 

2  Fe  Clo  +  2  H  Cl  +O  =  2  Fe  C13  +  H2  O  . 

On  the  other  hand,  if  an  excess  of  hydrochloric  acid  is  not  present, 
ferric  chloride  and  an  insoluble  basic  ferrous  chloride  are  formed :  — 

4  FeCl2  +  H20  +  O  =  2  Fe  j  ^H  +  2  FeCl3. 

Of  course,  the  usual  oxidizing  agents  (nitric  acid,  potassium  chlorate, 
and  hydrochloric  acid,  chlorine,  bromine,  etc.)  instantly  change 
ferrous  chloride  into  ferric  chloride.  Ferrous  chloride,  like  the  chlo- 
rides of  the  majority  of  other  divalent  elements,  forms  double  salts 
with  the  chlorides  of  the  alkalies,  as  well  as  with  the  chlorides  of  a 
number  of  other  metals. 

FERROUS  SULPHATE  (green  vitriol,  copperas),  Fe  S  O4  +  7  H2  O,  is  isomor- 
phous  with  the  vitriols  (see  page  417).  It  is  produced,  commercially, 
by  the  spontaneous  oxidation  of  iron  pyrites,  or  by  dissolving  iron  in 
dilute  sulphuric  acid  (see  page  32).  The  salt  loses  six  molecules  of 
water  at  100°,  and  is  completely  dehydrated  at  300°;  it  then  forms  a 
white  powder.  Like  the  other  vitriols,  ferrous  sulphate  forms  double 
salts  with  the  sulphates  of  the  alkali  metals ;  the  latter  contain  six 


480  FERROUS    COMPOUNDS. 

molecules  of  water  of  crystallization,  and  have  the  general  formula, 
Fe  SO4  ,  M2  SO  +  6  H2  O;  they  are  not  as  easily  oxidized  as  the  pure 
sulphate  of  iron.  When  exposed  to  moist  air,  ferrous  sulphate  is 
oxidized  to  a  mixture  of  ferric  sulphate  and  ferric  hydroxide:  *  — 

6  Fe  SO4  +  3  H2  O  +  3  O  =  2  Fe2  (SO4)3  +  2  Fe  (OH)3  , 

while,  in  the  presence  of  sulphuric  acid,  ferric  sulphate  alone  is  pro- 
duced :  — 

2  Fe  SO4  +  H2  SO4  +  O  =  Fe2  (SO4)3  +  H2  O. 

The  usual  laboratory  oxidizing  agents  have  the  same  effect.  Ferrous 
sulphate  absorbs  nitric  oxide  (NO),  forming  an  unstable  chemical 
compound  with  the  latter.  The  solution  is  dark  brown  in  color,  and 
parts  with  the  dissolved  gas  when  heated,  t 

FERROUS  SULPHIDE,  Fe  S,  sometimes  occurs  as  a  mineral.  It  is  formed 
by  heating  together  iron  and  sulphur  or  by  precipitation,  as  a  black 
powder,  when  the  solution  of  an  alkaline  sulphide  is  added  to  a  solu- 
tion of  a  ferrous  salt:  — 

Fe  S04  +  (NH4)2  S  =  Fe  S  +  (NH4)2  SO4  . 

Ferrous  sulphide  belongs  to  the  class  of  sulphides  which  are  dissolved 
by  dilute  acids  ;  it  is,  therefore,  not  precipitated  by  hydrogen  sulphide 
in  acid  solution  (see  page  100).  When  an  alkaline  sulphide  is  added 
to  a  solution  containing  a  ferric  salt,  not  ferric  sulphide,  but  ferrous 
sulphide  mixed  with  sulphur,  is  precipitated,  this  phenomenon  being 
due  to  the  instability  of  ferric  sulphide  in  aqueous  solution;  the  change 
may  be  represented  by  the  following  equations  :  — 


1.  Fe2S3  = 

2.  Fe2(SO4)3  +  3Na2S  =  2  FeS  +  S  +  3Na2SO4. 

FERROUS  CARBONATE,  Fe  CO3,  occurs  as  the  mineral  siderite,  isomorphous 
with  calcite.  It  can  be  formed  in  the  laboratory  by  adding  a  soluble 
carbonate  to  a  solution  containing  a  ferrous  salt  :  — 

Fe  S04  +  Na2  CO3  =  Na2  SO4  +  Fe  CO3  . 
Soluble.  Soluble.      Insoluble. 

When  moist,  ferrous  carbonate  is  readily  oxidized  by  the  air,  ferric 
hydroxide  remaining,  for  ferric  hydroxide  (like  the  hydroxide  of  alu- 

*  Possibly  a  basic  sulphate. 

t  The  formation  of  this  solution  is  a  delicate  test  for  nitric  acid.     Nitric 
acid  oxidizes  ferrous  sulphate  as  follows  :  — 

GFeS04+2HN03  +  3H2S04=3Fe2(S04)3  +  4H20  +  2NO; 

the  nitric  oxide,  which  is  liberated,  colors  the  excess  of  ferrous  sulphate.  Of 
course,  this  test  may  also  be  used  for  detecting  the  presence  of  a  nitrate,  for  the 
latter,  with  sulphuric  acid,  forms  a  sulphate  and  free  nitric  acid. 


FERRIC   OXIDE  ;    HYDROXIDE.  481 

minium  and  chromic  hydroxide)  is  too  weakly  basic  to  form  a  carbon- 
ate (see  page  342).  Ferrous  carbonate  is  easily  decomposed  into 
ferrous  oxide  and  carbon  dioxide  by  heat. 

Ferric  oxide,  Fe2  03 ,  is  found  in  nature  as  hematite;  in  the 
laboratory  it  can  be  produced  by  heating  the  corresponding  hydrox- 
ide, Fe  (OH  )8 :  — 

2  Fe_(OH  )8  =  Fe2  03  +  3  H2  0. 

The  oxide  so  prepared  has  a  fine  red  color ;  it  is  known  as  rouge,  and 
is  used  as  a  polish  for  metals  and  glass.  Ferric  hydroxide  is  pre- 
cipitated by  adding  an  alkaline  hydroxide  or  ammonia  water  to  a 
solution  containing  a  ferric  salt :  — 

Fe2  ( S04)3  +  6  KOH  =  2  Fe  (OH  )3  +  3  K2  S04 ; 

the  ferric  hydroxide  so  formed  probably  loses  water  spontaneously, 
so  that  the  precipitate  is  really  the  metahydroxide,  Fe  0  (OH).  A 
number  of  the  ferric  hydroxides,  which  are  formed  by  the  separa- 
tion of  water  between  two  or  more  formula  weights  of  the  normal 
hydroxide,  occur  as  minerals  which  are  commercially  important; 
such  compounds  are  limonite,  (4  Fe  (OH  )3  =  Fe4  03  (OH  )6  -f-  3  H2  0, 
and  xanthosiderite  (bog  iron  ore),  Fe2  0  (OH  )4 ;  *  the  latter  isomor- 
phous  with  beauxite  (see  page  333).  Ferric  hydroxide  can  also  be 
obtained,  in  a  soluble  form,  by  dialyzing  an  extremely  dilute  solu- 
tion of  ferric  chloride  mixed  with  ammonium  carbonate  (see  pages 
305,  306)  ;  the  solution  has  properties  similar  to  those  of  dialyzed 
silicic  acid.  Ferric  hydroxide  is  distinguished  from  the  hydroxide 
of  aluminium  and  from  chromic  hydroxide  by  the  fact  that  it  is  in- 
soluble in  an  excess  of  caustic  alkali ;  nevertheless,  it  does  possess 
acidic  properties,  as  is  proved  by  the  existence  of  minerals  like 
franklinite,  for  the  latter  is  a  zinc  ferrite,  Zn  (  Fe  02  )2 ,  derived  from 
meta  ferric  hydroxide,  Fe  0  (OH  )  ;  |  f ranklinite  is  isomorphous 
with  spinell  (see  page  339),  the  latter  being  a  salt  of  meta  alumin- 
ium hydroxide,  Al  O  (OH ),  acting  as  an  acid.  Besides  being 
acquainted  with  zinc  ferrite,  we  know  that  magnetite  is,  in  all 
probability,  a  ferrous  salt  of  ferric  hydroxide,  for  it  has  the  formula 
Fe  ( Fe  02)2 ,  and  is  isomorphous  with  spinell ;  and,  lastly,  calcium 

*  2  Fe  (OH%  -  Fe2  O  (OH)4  +  H2  O. 

t  In  franklinite  a  part  of  the  zinc  is  replaced  by  ferrous  iron  and  by  man- 
.ganous  manganese. 


482  FERRIC   SALTS. 

and  barium  ferrites  can  be  produced  in  the  laboratory  by  heating 
ferric  oxide  with  calcium  or  barium  oxide  to  a  high  red  heat ;  the 
latter  salts  have  the  composition :  — 

FejO  FejO 

^  >  Ca  and  ^  >  Ba. 

'"'  T71          \      {J 


On  the  other  hand,  ferric  oxide  and  ferric  hydroxide  are  bases,  for 
they  dissolve  in  acids  to  form  ferric  salts.* 

FERRIC  CHLORIDE,  Fe  C13 ,  can  be  produced  by  dissolving  ferric  hydroxide 
in  hydrochloric  acid,  or  by  passing  chlorine  into  a  solution  of  ferrous 
chloride :  — 

Fe  C12  +  Cl  =  Fe  C13 . 

Upon  evaporation  in  the  cold,  crystals  of  the  composition  Fe  Cl3-f  3  H2  O 
are  deposited;  another  hydrate  of  the  formula  Fe  C13  +  6  H2  O  is  formed 
if  the  solution  is  allowed  to  stand  at  ordinary  temperatures ;  a  solution 
of  ferric  chloride  cannot  be  heated,  because  then  it  breaks  down  into 
hydrochloric  acid  and  an  insoluble  oxychloride ;  t  reducing  agents,  such 
as  hydrogen  sulphide,  zinc  and  hydrochloric  acid,  sulphurous  acid,  etc., 
change  ferric  chloride  into  ferrous  chloride.  Anhydrous  ferric  chloride 
is  prepared  by  passing  chlorine  over  powdered  iron ;  the  chloride  evap- 
orates at  the  temperature  of  boiling  sulphur  (448°),  and  its  vapor  den- 
sity, even  at  that  temperature,  is  too  low  for  a  substance  consisting  of 
molecules  having  the  formula  Fe2  C16 ,  the  specific  gravity  of  the  gas- 
eous chloride  rapidly  decreases  as  the  temperature  is  raised,  so  that 
no  doubt  can  exist  as  to  the  trivalence  of  iron  in  ferric  chloride,  the 
molecule  of  which,  in  a  gaseous  state,  is  expressed  by  the  formula 
FeCl8,  (see  page  336).}  Ferric  chloride  combines  with  other  chlo- 
rides to  form  double  salts  similar  to  those  produced  by  aluminium 
chloride  (see  page  337  ). 

FERRIC  SULPHATE,  Fe2  (SO4  )3  is  formed  by  oxidizing  ferrous  sulphate 
(see  page  340)  or  by  dissolving  ferric  hydroxide  in  sulphuric  acid;  it 
forms  alums  with  the  sulphates  of  the  alkali  metals  (see  page  340  ). 

Ferric  sulphide,  Fe2  S3 ,  analogous  to  ferric  oxide,  Fe2  03 ,  can  be 
produced  by  heating  ferrous  sulphide  and  sulphur  to  a  red  heat :  — 


*  Ferric  oxide,  like  the  oxides  of  aluminium  and  chromium,  is  not  dis- 
solved by  acids  after  it  has  been  heated  to  a  red  heat. 

t  If  the  solution  is  quite  dilute,  complete  decomposition  and  formation  of 
soluble  ferric  hydroxide  takes  place: — 

Fe  C13  +  3  H2  O  =  Fe  (OH)8  +  3  H  Cl. 

t  W.  Grunewald  and  V.  Meyer,  Berichte  der  Deutsch.  Chem.  Gesell.  ;21, 
687. 


FERROUS-FERRIC    OXIDE.  483 

it  has  the  character  of  an  acidic  anhydride,  and  forms  a  number  of 
salts  which  occur  as  natural  minerals  ;  an  example  of  these  com- 
pounds is  chalcopyrite  (copper  pyrites),  in  which  cuprous  sulphide 
is  the  base  :  — 

Cu2  S  +  Fe2  S3  =  Cu2  Fe2  S4  . 

These  sulphoferrites  are  analogous  to  the  ferrites  (examples  of 
the  latter  occur  in  the  spinell  group)  j  when  heated,  ferric  sulphide 
is  converted  into  the  compound  Fe3S4;  this  change  is  similar  to 
that  undergone  by  ferric  oxide. 

Ferrous-ferric  oxide,  Fe3  04  ,  occurs  as  magnetite,  a  mineral  which 
has  the  power  of  being  attracted  by  the  magnet.  The  oxide  is  pro- 
duced when  iron  is  burned  in  an  excess  of  oxygen,  or  when  steam  is 
passed  over  red-hot  iron  (see  page  31  )  ;  the  black  coating  which 
forms  on  iron  which  is  heated  to  a  high  red  heat  consists  of  a  mix- 
ture of  Fe8  04  and  Fe2  03  .  Ferrous-ferric  oxide  is  considered  to  be 
constituted  similarly  to  the  spinells  ;  its  structure  would,  therefore, 
be  represented  as  follows  :  — 

Ho 

>  Fe. 


The  sulphide,  Fe3  S4  ,  which,  in  formula,  corresponds  to  magnetite, 
is  produced  when  a  current  of  the  hydrogen  sulphide  is  passed  over 
red-hot  iron  :  — 

3  Fe  +  4  H2  S  =  Fe3  S4  +  8  H. 

Like  the  oxide,  it  has  magnetic  properties. 

Ferric  acid  is  not  known  in  the  free  state  ;  the  ferrates,  which  in 
formula  are  analogous  to  the  sulphates,  chromates,  and  manganates, 
however,  exist  in  limited  numbers  ;  the  best  known  of  these  is  potas- 
sium ferrate,  K2  Fe  04  ;  the  latter  is  produced,  just  as  is  potassium 
manganate,  by  fusing  the  metal  with  potassium  nitrate  ;  the  same 
salt  is  also  formed  by  passing  chlorine  into  a  solution  of  potassium 
hydroxide  which  contains  ferric  hydroxide  in  suspension.*  Potas- 

*  The  oxidizing  agent  is  the  potassium  hypochlorite  which  is  produced 
(see  page  121)  ;  the  solution  of  ferrate  of  potassium  has  a  violet  color.  Ordi- 
nary potassium  hydroxide  not  infrequently  contains  some  ferric  hydroxide; 
and  the  formation  of  potassium  ferrate,  with  its  violet  color,  is  sure  to  puzzle 
the  student  when  he  is  preparing  potassium  hypochlorite. 


484  COBALTOUS  COMPOUNDS. 

slum  ferrate,  as  well  as  the  other  ferrates,  is  unstable ;  when  allowed 
to  stand  it  breaks  down,  giving  off  oxygen  and  leaving  ferric 
hydroxide. 

The  disulphide  of  iron,  Fe  S2 ,  has  no  analogon  among  the  oxygen 
compounds  of  the  metal ;  it  is,  however,  extremely  important,  be- 
cause it  occurs  so  widely  distributed  as  the  mineral  iron  pyrites. 
The  disulphide  is  dimorphous  ;  as  pyrites  it  crystallizes  in  the  regu- 
lar system,  and  as  markasite  it  is  rhombic ;  it  has  a  metallic  lustre, 
and,  on  a  superficial  examination,  has  somewhat  the  appearance  of 
gold. 

The  cyanides  of  iron,  the  ferro  and  ferricyanides,  as  well  as  the 
corresponding  acids,  were  discussed  on  page  296. 

Compounds  of  Cobalt.  Cobalt  forms  two  series  of  compounds, 
cobaltous  compounds,  derived  from  cobaltous  oxide,  Co  0,  and  cobaltic 
compounds,  derived  from  cobaltic  oxide,  Co2  03 ;  the  cobaltous  salts 
are  the  ones  which  are  most  frequently  met  with,  cobaltic  oxide 
having  only  very  weakly  basic  properties. 

Cobaltous  oxide,  Co  0,  can  be  prepared  by  decomposing  the  corre- 
sponding hydroxide,  Co(OH)2,  by  heat,  air  being  excluded;  the 
hydroxide  is  precipitated  by  adding  an  excess  of  caustic  alkali  to 
the  solution  of  a  cobaltous  salt :  — 

Co  (  N03  )2  +  2  KOH  =  Co  (OH  )2  +  2  KN08 . 

Cobaltous  hydroxide  is  rose  colored,  the  oxide  is  olive  green ;  both 
the  oxide  and  hydroxide  are  strong  bases,  dissolving  in  acids  to 
form  cobaltous  salts. ' 

COBALTOUS  CHLORIDE,  Co  C12,  can  be  formed  by  dissolving  the  carbonate  or 
hydroxide  in  hydrochloric  acid;  when  the  solution  is  evaporated,  red 
crystals  of  the  composition  Co  C12  +  6  H2  O  separate ;  when  heated 
to  100°  the  salt  becomes  violet  and  has  the  composition  Co  C12  + 
H2O;  at  110°-120°  it  loses  all  its  water  of  crystallization,  and  is  con- 
verted into  the  anhydrous  chloride,  Co  C12 ,  which  is  blue ;  the  same 
rule  appertains  to  all  cobalt  salts;  when  hydrated  they  are  red  or 
rose-colored,  when  anhydrous  they  are  blue.* 

*  This  property  of  cobalt  salts  caused  the  use  of  those  substances  as  far 
back  as  1757  for  the  preparation  of  sympathetic  ink;  this  consists  of  cobalt 
chloride  in  solution,  and  is,  therefore,  pale  red,  and  almost  invisible  on  the 
paper.  When  the  paper  is  heated,  however,  the  salt  becomes  anhydrous,  and 
the  writing  appears  in  plain,  blue  characters. 


COBALTIC   COMPOUNDS.  485 

COBALTOUS  NITRATE,  Co  (NO3)2  +  6  H2  O,  is  the  most  common  cobalt  salt; 
it  loses  water  at  100°,  and  at  a  higher  temperature  decomposes,  giving 
off  nitrogen  dioxide  and  leaving  cobaltic  oxide. 

A  COBALT  GLASS,  formed  by  fusing  silicon  dioxide,  potassium  carbonate, 
and  some  salt  of  cobalt,  has  an  intensely  blue  color;  when  finely  ground, 
it  is  known  as  smalt ;  the  glass  has  approximately  the  same  composition 
as  ordinary  potash-glass  (see  page  421)  with  the  exception  that  the  cal- 
cium is  replaced  by  cobalt.  Fused  borax  will  dissolve  cobalt  oxide,  or 
salts  of  cobalt,  leaving  an  intensely  blue,  glass-like  substance  which 
contains  cobaltous  metaborate ;  the  same  is  true  of  sodium  metaphos- 
phate.  The  formation  of  these  two  compounds  is  a  ready  means  for 
detecting  the  presence  of  cobalt.  When  heated  with  aluminium 
oxide,  cobalt  compounds  form  a  blue  cobalt  aluminate  which  is  known 
as  Thenard's  blue. 

COBALTOUS  SULPHIDE,  Co  S,  is  a  black  precipitate,  formed  by  adding  an 
alkaline  sulphide  to  a  solution  containing  a  cobalt  salt :  — 

Co  (N03)2  +  (NH4)2  S  =  Co  S  +  2  NH4  NO3; 

it  differs  from  ferrous  sulphide  by  not  being  soluble  in  very  dilute 
acids. 

Cobaltic  oxide,  Co2  08 ,  has  very  weakly  basic  properties  ;  its  salts, 
which  are  formed  by  dissolving  the  oxide  in  cold  acids,  are  easily 
decomposed,  and  are  tolerably  stable  only  in  solution ;  a  number  of 
double  salts,  derived  from  a  cobaltic  salt  united  with  a  salt  of  some 
other  base,  are  known,  which  are  more  stable  than  the  pure  cobaltic 
salts. 

COBALTIC  NITRITE  combines  with  potassium  nitrite  to  form  a  double  salt 
of  the  composition  Co(NO2)3,  3KNO2,  which  is  of  importance  be- 
cause it  is  insoluble  in  water  and  in  cold,  dilute  acids ;  its  precipitation 
may,  therefore,  be  used  in  separating  cobalt  from  nickel.  It  is  produced 
by  adding  potassium  nitrite  to  a  slightly  acid -solution  of  a  cobalt  salt; 
the  nitrous  acid,  which  is  liberated  from  the  potassium  nitrite,  then 
oxidizes  the  cobaltous  compound  to  a  cobaltic  one,  after  which  reaction 
the  double  nitrite  of  cobalt  and  potassium  is  precipitated. 

A  solution  of  cobaltous  chloride,  when  in  the  presence  of  am- 
monia,* and  exposed  to  the  air,  is  slowly  oxidized ;  by  this  means 

*  The  addition  of  an  excess  of  ammonia  to  cobaltous  salts  does  not  produce 
a  precipitate  of  hydroxide ;  the  same,  it  will  be  remembered,  is  true  of  cupric 
salts. 


486  COBALT   AMINES. 

there  is  then  formed  a  peculiar  series  of  compounds  composed  of 
cobaltic  chloride  united  with  a  varying  number  pf  molecules  of  am- 
monia. These  compounds  are  known  as  cobalt  amines;  the  structural 
composition  of  the  latter,  despite  the  extended  amount  of  work 
which  has  been  done  upon  them,  is,  as  yet,  not  understood ;  analysis 
shows  us  that  one  atom  of  cobalt  can  unite  with  three,  four,  five,  and 
six  molecules  of  ammonia  to  form  different  radicles  which  can  play 
the  part  of  individual  trivalent  bases  ;  these  bases,  themselves,  exist 
only  in  aqueous  solution,  the  salts  derived  from  them  are,  however, 
stable ;  the  ammonia  in  these  compounds  cannot,  therefore,  be  com- 
pared to  water  of  crystallization,  as  it  is  in  the  case  of  the  ammonia 
compounds  of  copper  sulphate  (see  page  403).  Only  a  few  of  the 
simpler  connections  can  be  traced  here ;  for  a  more  extended  review 
of  this  complicated  subject  the  pupil  must  refer  to  a  larger  text- 
book.* 

r  1.  Dichrocobaltic  chloride,  Co(NH3)3  C13  +  H2  O. 

2.  Praseocobaltic  chloride,  Co  (NH3)4  C13  +  H2  O. 

From  solution  of  Co  C12  J  .  Roseocobaltic  chioride,  Co  (NH3)5  C13  +  H2  O. 

in  air  are  produced  •  (  Purpureocobaltic  chloride,  Co  (XH,)5  C13  t. 

1 4.  Luteocobaltic  chloride,  Co  (NH3)6  C13 . 

The  radicles  which  assume  metallic  functions  in  this  series  of  salts 
may  be  compared  with  metallic  cobalt ;  for  from  each,  in  addition  to 
the  chlorides,  a  number  of  other  salts,  such  as  the  nitrate  and  sul- 
phate, are  derived.  The  relationship  is  demonstrated  by  the  follow- 
ing table  :  — 

Co  Cl  3,  cobaltic  chloride;  Co  (NO3)3  ,  nitrate;  Co2  (SO4)3  ,  sulphate;     ' 

Co(NH3)3Cl3,  dichro  chloride ;  Co(NH3)3(NO3)3  ,  nitrate;  [Co(NH3)3]2  (SO4)3  ,  sulphate; 
Co(NH3)4Cl3,  praseo  chloride;  Co(NH3)4(NO3)3  ,  nitrate;  [Co(NH3)4]2  (SO'4)3 ,  sulphate; 
Co(NH,).  C13,  roseo  chloride;  Co(NH3)6(NO3)3  ,  nitrate;  [Co(NH3)5]2  (SO4)3  ,  sulphate; 
Co(NH3)6Cl3,  luteo  chloride;  Co(NH3)6(NO3)3  ,  nitrate;  [Co  (NH3)6]2(SO4)3  ,  sulphate. 

The  cyanides  of  cobalt  correspond  to  those  of  iron  (see  page 
296). 

The  compounds  of  nickel  are  almost  exclusively  derived  from 
nickelous  oxide.  The  latter  can  readily  be  produced  by  heating  the 
hydroxide,  Ni  (OH  )2 ,  which  is  precipitated  as  a  grass-green  sub- 
stance on  adding  a  soluble  hydroxide  to  a  solution  containing  a  salt 

*  Ladenburg;  Handworterbuch  der  Chemie  ;  volume  5. 

t  Roseocobaltic  chloride  is  formed  from  cobaltic  chloride  and  concentrated 
ammonia  in  the  cold;  purpureocobaltic  chloride,  by  boiling  the  same  with  con- 
centrated hydrochloric  acid.  The  difference  is  in  the  water  of  crystallization. 


NICKEL;  COMPOUNDS  OF.  487 

of  nickel.*     Nickel  salts,  when  hydrated,  are  green;  when  anhy- 
drous, yellow. 

NICKELOUS  CHLORIDE,  Ni  C12  +  6  H2  O,  is  formed  by  dissolving  the 
oxide  or  hydroxide  in  hydrochloric  acid.  Nickelous  sulphate,  pro- 
duced by  substituting  sulphuric  for  hydrochloric  acid,  is  isomorphous 
with  the  vitriols  (see  page  417).  Nickelous  sulphide  is  precipitated 
as  a  black  powder  when  a  solution  of  an  alkaline  sulphide  is  added  to 
a  solution  containing  a  nickelous  salt:  — 


Ni  S04  +  (NH4)2  S  =  Ni  S  +  (NH4)2  SO4  . 
Nickelous  sulphide  is  insoluble  in  dilute  acids. 

The  cyanides  of  nickel  do  not  correspond  to  those  of  iron  and  of 
cobalt,  with  the  exception  of  nickelous  cyanide.  The  latter,  with  a 
formula  of  Ni  (CN  )2  ,  is  precipitated  when  potassium  cyanide  is 
added  to  a  solution  containing  a  salt  of  nickel  ;  on  addition  of  an 
excess  of  the  reagent,  it  forms  a  double  cyanide  of  the  formula 
Ni  (CN  )2  ,  2  KCN,  which,  therefore,  does  not  correspond  to  the 
potassium  salts  of  ferrocyanic  and  cobaltocyanic  acids  :  — 

Ni  (CN  )2  ,  2  KCN.  K4  Fe  (CN  )6  .  K4  Co  (CN  )6  . 

Double  salt  of  nickel.        Potassium  ferrocyanide.        Potassium  cobaltocyanide. 

Potassium  ferrocyanide  and  cobaltocyanide  can  be  readily  oxidized 
to  the  cobaltic  and  ferric  compounds,  K3  Fe  (CN  )6  and  K3  Co  (CN  )6  , 
which  are  salts  of  cobalticyanic  acid  and  ferricyanic  acid  respec- 
tively ;  but  the  nickel  cyanide  is  not  capable  of  oxidation.  Nickel 
can  be  precipitated  from  its  double  cyanide  by  the  usual  reagents, 
such  as  ammonium  sulphide  ;  but  the  iron  or  cobalt  cannot  be 
separated  by  any  such  means,  when  either  has  entered  into  the 
formation  of  the  peculiar  acids  from  which  its  double  cyanides 
are  derived.  In  the  behavior  of  its  cyanides,  in  the  isomorphism  of 
its  sulphate  with  the  vitriols,  and  in  a  number  of  other  respects, 
nickel  resembles  the  first  element  in  the  second  half  of  the  long 
period,  i.e.,  copper. 

The  relationship  between  the  formulae  of  the  compounds  of  iron, 
cobalt,  nickel,  chromium,  manganese,  and  sulphur  are  given  in  the 
following  table  :  — 

*  Excess  of  ammonia  water  dissolves  the  hydroxide  ;  this  is  similar  to  the 
action  of  cobaltous  hydroxide  and  cupric  hydroxide. 


488  IRON  ;   COBALT  ;    NICKEL  ;    TABLE  OF. 

OXIDES   WHICH   ARE   BASIC,    AND    SALTS   DERIVED    FROM   THEM. 


-ous 

OXIDES. 

-Oils    SAI.'IS. 

-ic 

OXIDES. 

-1C   SALTS. 

CrO 
MnO 
FeO 
CoO 
NiO 

CrCla 
MnCl, 
FeCl2 
CoCl2 
NiCl2 

chlorides 

CrSO4* 
MnSO4 
FeSO4 
CoSO4 
NiSO4 

sulphates 

Cr203$ 
Mn2O3 
Fe203* 
Co203 
Ni203 

CrCl3 
Mn  Cl3t 
Fe  Cl3t 
CoCl3t 

Cr2fS04)3 
Mn2(S04)3| 
Fe2(S04)3t 

Cr(NO3)3 

Mn(N03)2 
Fe(N03)2 
Co(N03)2 
Ni(N03), 
nitrates 

Fe(N03)3t 

chlorides 

sulphates 

nitrates 

*  Compounds  marked  *  are  unstable,  very  readily  oxidized ;  those  marked  t 
readily  change  to  -ous  compounds.  Ni2  O3  is  known  as  its  hydroxide  Ni  (OH)3 , 
formed  by  oxidizing  a  nickelous  salt  with  potassium  hypochlorite  or  hypobro- 
mite;  it  forms  no  salts. 

The  -ous  compounds  are  converted  into  -ic  compounds  by  oxidation,  the 
-ic  compounds  into  -ous  compounds  by  reduction. 

J  Both  basic  and  acidic,  Cr2  O3  dissolves  in  caustic  alkalies  to  form  chro- 
mites;  Fe2  O3  forms  some  compounds  which  are  analogous  to  the  chroinites  and 
aluminates;  these  salts  are  derived  from  a  hydroxide  MO  (OH). 


OXIDES   WHICH   ARE   COMPOSED    OF    THE    TWO    GIVEN    ABOVE. 

-ous-ic-   OXIDES. 

Cr304,  Mn304,  Fe3O4,  Co3  O4. 

Of  the  elements  in  this  table,  manganese  alone  forms  a  dioxide,  Mn  O2;  how- 
ever, sulphides  or  arsenides  of  the  other  elements,  analogous  in  structure, 
are  known. 


OXIDES   WHICH   ARE   ACIDIC    ONLY. 


a.   Acids  and  salts  of  the  general  formula  H2  XO4;  type  H2  SO4 


ANHYDRIDES. 


CrO3 


M2  Cr  O4  ,  chromates ; 
M2  Mn  O4  ,  manganates ; 
M2  Fe  O4  ,  ferrates. 


b.   Acids  and  salts  of  the  general  formula  HXO4;  type  H  Cl  O4 


ANHYDRIDE. 

SALTS. 

Mn.2  O7 

M  Mn  O4  ,  permanganates. 

ELEMENTS    OF    PLATINUM    GROUP.  489 


CHAPTER   LX. 

THE    REMAINING    ELEMENTS    OF    THE    EIGHTH   FAMILY. 
(THE    PLATINUM    GROUP.) 

Ruthenium  ;  symbol,  Ru  ;  atomic  weight,  101.6  ; 
Rhodium  ;  symbol,  Rh  ;  atomic  weight,  103  ; 
Palladium  ;  symbol,  Pd  ;  atomic  weight,  106.6 ; 
Osmium  ;  symbol,  Os ;  atomic  weight,  190.8  ; 
Iridium  ;  symbol,  Ir  ;  atomic  weight,  193.1 ; 
Platinum ;  symbol,  Pt ;  atomic  weight,  195. 

THE  six  elements  named  above  are  all  extremely  rare,  and,  with 
the  exception  of  platinum,  of  little  practical  importance ;  they  fall 
into  two  groups  represented  by  the  horizontal  lines  in  the  following 
table,  and  also  into  three  groups,  represented  by  the  vertical  col- 
umns, as  follows :  —  Ku>  Kh  pd 

Atomic  weight,  101.6         103.         106.6 

Os.  Ir.  Pt. 

Atomic  weight,  190.8         193.1       195. 

In  passing  from  left  to  right  in  each  of  the  two  horizontal  lines, 
we  find  a  gradation  in  properties  similar  to  that  observed  in  the  group 
composed  of  iron,  cobalt,  and  nickel  ;  thus,  ruthenium  and  osmium 
are  able  to  form  higher  oxides  than  the  remaining  four,  just  as  iron 
is  able  to  form  a  higher  oxide  than  either  cobalt  or  nickel ;  and  the 
same  distinction  in  the  formulae  of  the  double  cyanides  which  was 
observed  as  existing  between  those  of  iron  and  cobalt  on  the  one 
hand  and  nickel  on  the  other,  is  observed  between  ruthenium  — 
rhodium,  osmium  —  iridium,  and  palladium  —  platinum  ;  this  dis- 
tinction is  made  clear  by  the  following  table :  — 


K4Fe(CN)6 
K4  Co  (CN)6 
K4Ru(CN)6 
K4Os(CN)6 

K4Fe(CN)6 
K4Co(CN)6 
K4Rh(CN)6 
K4  Ir  (CN)6  ^ 

Fe,  Co;  Rh,  Ir. 

Ni(CN)2,  2KCN 
Pd(CN)2,  2KCN 
Pt(CN)2,  2KCN 

Fe,  Co;  Ru,  Os. 

Ni;Pd,  Pt. 

490 


ELEMENTS  OF  PLATINUM  GKOUP. 


The  difference  existing  between  the  oxides  may  also,  perhaps, 
be  best  given  in  the  form  of  a  table.  The  highest  oxides  belong 
to  ruthenium  and  osmium;  indeed,  these  are  the  only  two  ele- 
ments which  are  able  to  form  oxides  of  the  formula  M04*  in 
which  the  valence  of  the  element  is  presumably  eight,  provided 
we  regard  oxygen  as  having  a  valence  of  two.  The  highest  val- 
ence toward  oxygen  displayed  by  any  element,  therefore,  appears 
in  the  eighth  family.  In  the  following  table  the  formulae  of  the 
oxides  of  the  six  elements  under  discussion  are  compared  with 
those  of  manganese,  the  element  belonging  to  the  preceding 
(VII.)  family;  as  will  be  noticed,  there  is  a  great  resemblance, 
although  the  highest  valence  of  manganese  toward  oxygen  is  only 
seven : — 


Mn. 

Ru  and  Os. 

Rh  and  Ir. 

Pd  and  Ft. 

Mn2O7 
(Mn03)J 
MnO2 
Mn203 
MnO 

RuO4;   OsO4 
(Ru.207)t;- 
(BuOa)t;  (Os03)t 
RuO2;  OsO2 
Ru203;  Os203 
RuO;  OsO 



(IrOg)§ 
RhO2;    IrO2 
Rh203;  Ir203 
RhO;     IrO 

Pd02;  Pt02 
PdO;  PtO 

t  The  anhydride  is  not  known;  the  salts,  the  per-ruthenites,  M  Ru  O4 , 
corresponding  to  the  permanganates,  are  known. 

|  The  anhydrides  are  not  known,  but  the  salts,  manganates,  ruthenites, 
and  osmites  are  known ;  their  general  formula  is  M2  X  O4 ,  corresponding  to  the 
.sulphates.  An  oxide  Pd2  O  is  also  known. 

§  One  salt,  K2  Ir  O4,  derived  from  such  an  anhydride,  has  been  described. 

In  the  periodic  system,  the  element  following  palladium  is  sil- 
ver, and  the  one  following  platinum  is  gold ;  as  a  consequence,  we 
find  considerable  resemblance  between  palladium  and  silver  (for 
instance,  in  the  formation  of  the  oxide  Pd2  0),  and  between  plati- 
num and  gold. 

The  metals  of  the  platinum  group  all  possess  a  grayish-white 
color  and  brilliant  metallic  lustre.  They  fuse  at  a  high  white  heat, 
the  melting  points  decreasing  in  the  two  groups  from  ruthenium  to 


*  An  oxide  of  sulphur,  S  O4,  is  also  mentioned. 


ELEMENTS   OF   PLATINUM   GROUP  ;   OCCURRENCE.         491 

palladium,  and  from  osmium  to  platinum ;  their  specific  gravities 
and  atomic  volumes  are  as  follows  :  — 

Eu.  Eh.  Pd. 

Specific  gravities,        12.261  12.1  11.4 
Atomic  volumes,           8.06               8.5  9.3 

Os.  Ir.  Pt. 

Specific  gravities,        22.4  22.42  21.46 
Atomic  volumes,           8.55               8.6  9.08. 

The  elements  of  the  platinum  family  are,  therefore,  at  the  mini- 
mum of  the  curves  of  atomic  volumes  belonging  to  their  respective 
periods ;  the  next  following  elements,  with  larger  atomic  weights, 
have  larger  atomic  volumes.  The  platinum  metals,  are  (like 
iron,  cobalt,  and  nickel)  malleable  and  ductile,  and  form  colored 
salts.  All  of  the  elements  which  have  small  atomic  volumes  and 
high  specific  gravities  display  several  oxides  and  series  of  com- 
pounds ;  in  none  of  these,  however,  is  the  chemical  character  of  any 
such  element  with  small  atomic  volume  so  pronounced  as  to  over- 
shadow the  character  of  the  other  elements  in  combination  ;  as  has 
been  previously  remarked,  therefore,  the  crowding  of  a  large  amount 
of  matter  into  a  small  space  seems  to  be  unfavorable  for  manifesta- 
tion of  very  decided  metallic  or  not-metallic  properties.  All  of  the 
elements  belonging  to  the  eighth  family  possess  the  power  of  con- 
densing and  transmitting  hydrogen  (see  page  34). 

The  members  of  the  platinum  family  all  occur  as  the  uncom- 
bined  metals,  they  are  never,  however,  pure  in  their  natural  condi- 
tion ;  platinum  may  be  found  with  only  a  slight  admixture  of  iron, 
but,  as  a  general  rule,  alloys  of  the  various  metals  forming  the  plat- 
inum group  (which  may  also  contain  silver  and  gold)  occur. 

The  metals  of  this  group  are  very  difficult  to  dissolve  in  acids  ; 
ruthenium  and  iridium  are  not  attacked  by  any  acid ;  platinum  is 
changed  only  by  aqua  regia ;  osmium  is  dissolved  by  nitric  acid  as 
well ;  while  palladium  is  soluble  both  in  nitric  and  hydrochloric 
acids  ;  all  of  the  elements  are  very  easily  reduced  and  isolated  from 
their  oxides  or  halides  ;  they  all  manifest  a  great  tendency  to  unite 
with  chlorine.  In  their  chlorine  compounds,  ruthenium,  rhodium, 
and  palladium  resemble  iron,  cobalt,  and  nickel,  for  they  form 
chlorides  with  the  formulae  M  CL  and  M  C13 ;  on  the  other  hand, 
osium,  iridium,  and  platinum  have  their  most  stable  chlorides  de- 


492  PLATINIC    CHLORIDE. 

rived  from  a  tetravalent  metal,  with  the  general  formula  M  C14 ;  the 
latter  compounds  have  a  great  tendency  to  produce  double  salts 
with  the  chlorides  of  the  alkali  metals ;  these  chlorides  have  the 
composition  M2  R  C16  ,*  and  may  be  regarded  as  derived  from  an  acid 
M2  E,  C16 ,  in  which  R  C14  acts  as  an  acidic  anhydride  ;  the  structure 
of  these  acids  would,  therefore,  be  analogous  to  that  of  fluosilicic 
acid  (see  page  303). 

The  higher  oxides  of  this  group  (mentioned  in  the  table,  page 
490)  are  acidic  in  their  character,  the  lower  ones  are,  at  best,  weakly 
basic.  The  two  tetroxides,  Ru  04  and  Os  04  ,f  are  quite  volatile,  and 
are  produced  by  powerful  oxidation  of  the  respective  metals ;  they 
have  a  most  peculiar,  penetrating  odor,  which  has  been  compared  to 
that  of  ozone  ;  $  the  vapor  densities  correspond  to  molecules  of  the 
formula  R  04 . 

Derived  from  a  next  lower,  hypothetical,  oxide,  in  which  ruthe- 
nium would  be  heptavalent,  we  have  one  salt,  K  Eu  04 ;  osmium  does 
not  furnish  a  parallel  compound.  Next  to  this  we  have  salts  of  two 
hypothetical  acids,  H2  Ru  04  and  H2  Os  04 ,  which  correspond  to 
manganic,  chromic,  and  sulphuric  acids ;  the  anhydrides,  were  they 
known,  wrould  have  the  formulae  Ru  03  and  Os  03 ;  the  salts  are 
termed  ruthenites  and  osmites  ;  only  the  potassium  salt  of  ruthenious 
acid  is  known,  but  the  potassium,  sodium,  and  barium  salts  of  osmous 
acid  have  been  isolated. 

The  dioxide,  M02 ,  is  the  highest  oxide  which  is  common  to  all 
of  the  platinum  metals ;  it  is  also  the  most  important  of  the  oxides 
of  this  group,  because  the  principal  salts  are  derived  from  it,  the 
chief  chlorides  of  this  family  having  the  formula  M  C14 .  The  dioxide 
is  basic  in  its  character,  although  that  of  platinum  is  also  slightly 
acidic,  for  it  dissolves  in  concentrated  alkalies  to  form  platinates. 
The  only  chloride  which  need  be  mentioned  in  detail  is  that  of 
platinum ;  its  characteristics  may  serve  as  a  type  for  all  of  the  rest. 

PLATINIC  CHLORIDE,  Pt  C14  +  5  H2  O,  is  prepared  by  adding  silver  nitrate 
solution  to  the  substance  which  is  ordinarily  termed  platinic  chloride, 
but  which  is,  in  reality,  an  acid  of  the  formula  H2Pt  C16;  the  first  reac- 
tion is  as  follows :  — 

*  M  represents  a  univalent  alkali  metal,  R  either  osmium,  iridium,  or 
platinum. 

t  The  latter  is  sometimes  called  osmic  acid. 

t  The  odor  is  almost  exactly  like  that  of  chinon. 


CHLORPLATLNTC   ACID.  493 

H.2  Pt  C16  +  2  Ag  NO3  =  Ag2  Pt  Cle  +  2  HNO3; 

the  silver  salt  of  chlorplatinic  acid,  which  is  precipitated,  is  decomposed 
by  boiling  water  as  follows :  — 

Ag2  Pt  C16  =  2  Ag  01  +  Pt  C14; 

the  insoluble  silver  chloride  is  filtered,  and  the  solution  evaporated  to 
crystallization. 

Chlorplatinic  acid,  H2  Pt  C16  +  6  H2  0,  is  obtained  by  dissolving 
platinum  in  aqua  regia  and  cautiously  evaporating ;  it  then  forms 
ochre-colored,  deliquescent  prisms.  The  chlorplatinates  are  of  im- 
portance in  chemical  analysis,  because  ammonium  and  potassium 
chlorplatinates  are  nearly  insoluble  in  water,  and  entirely  insoluble 
in  alcohol.  Potassium  chlorplatinate  is  precipitated  as  a  yellow, 
crystalline  substance,  when  potassium  chloride,  or  any  other  potas- 
sium salt,  is  added  to  a  solution  of  chlorplatinic  acid ;  its  formula  is 
K2  Pt  C16 .  Similar  results  are  obtained  with  ammonium  salts.  Am- 
monium chlorplatinate,  when  heated,  decomposes  entirely,  leaving 
platinum  in  a  very  finely  divided,  spongy  condition  in  which  it  is 
known  as  "  spongy  platinum."  Spongy  platinum  in  a  marked  degree 
possesses  the  power  of  occluding  gases,  and  its  use  has  been  mentioned 
in  several  of  the  preceding  chapters  (see,  for  instance,  the  prepara- 
tion of  sulphur  trioxide,  page  144).*  The  use  of  spongy  platinum 
depends  on  the  fact  that  occluded  gases  are  chemically  much  more 
active  than  they  are  in  the  ordinary  condition  (see  page  33).  Pla- 
tinic  bromide  or  iodide  forms  acids  similar  to  chlorplatinic  acid.  The 
structure  of  these  acids  is  similar  to  that  of  fluosilicic  acid,  H2  Si  F6 , 
which  is  also  derived  from  a  tetrahalide,  Si  F4 ;  so  that  in  these 
compounds,  also,  it  is  best  to  accept  a  theory  of  the  divalence  of  the 
halogen  atoms  (see  pages  303  and  337). 

Platinic  chloride  is  able  to  unite  with  ammonia  to  form  platinic- 
amines,  in  which  the  chemical  behavior  of  ammonia  is  identical  with 
that  of  the  same  substance  in  the  cobaltamine  salts.  Platinic  chlo- 

*  "  Platinized  asbestos  "  is  formed  by  dipping  asbestos  in  platinum  chlo- 
ride solution  and  then  in  one  of  ammonium  chloride;  ammonium  chlorplati- 
nate is  thus  precipitated  on  the  fibres,  and  can  be  changed  to  spongy  platinum 
by  heating.  Palladium  asbestos  may  be  prepared  in  the  same  way.  "Plati- 
num black"  is  very  finely  divided  platinum,  formed  by  reducing  a  solution  of 
platinum  chloride,  to  which  potassium  hydroxide  has  been  added,  by  means 
of  alcohol. 


494  CIILOR-PLATINAMINES. 

ride  forms  a  number  of  these  compounds,  only  two  of  which  is  it 
necessary  to  mention  here. 

Platinic  chloride  can  unite  with  two  molecules  of  ammonia  to 
form  chlor-platinamine  chloride,  which  has  the  formula  C12  Pt  (  NH3  )2 
C12 :  the  radicle  Pt  (  NH3  )2  being  tetravalent  and  uniting  with  four 
atoms  of  chlorine,  just  as  one  atom  of  platinum  does  in  platinic 
chloride :  — 

Pt  C14 ,  and  Pt  (  NH3  )2  C14 , 

Platinum  chloride,  and  chlor-platinamine  chloride. 

However,  in  the  latter  compound,  two  atoms  of  chlorine  can  be 
replaced  by  other  acid  groups  (-N03,  or  =S04)  very  readily, 
while  the  other  two  are  much  less  reactive;  this  compound  may 
therefore  also  be  considered  as  the  chloride  of  a  divalent  radicle 
Cl2Pt  (ISTH3)2  =.  Because  of  the  difficulty  with  which  they  are 
replaced,  it  is  supposed  that  the  two  less  reactive  chlorine  atoms  are 
attached  to  platinum,  while  the  other  two  are  united  to  nitrogen  in 
the  same  way  as  one  atom  of  chlorine  is  in  ammonium  chloride. 
The  structure  of  the  ammonium  compound  is  therefore  better  repre- 
sented by  the  following  formula :  — 

ri  p    (  NH3  Cl 

^12  r  l>  •<    T^-TT      pi 

and  the  nitrate  and  sulphate  by  the  following :  — 
N0<     UKlCLI 


This  theory  is  borne  out  by  the  fact  that  platinous  chloride,  Pt  C12 , 
can  form  a  similar  compound  with  ammonia,  with  the  difference 
that  in  this  case  two  chlorine  atoms,  attached  to  platinum,  are 
missing ;  this  compound  therefore  has  the  structure  :  — 

(NH3C1 

*  JNH3C1. 

In  the  second  series  of  platinic  amines,  which  will  be  described, 
there  are  four  molecules  of  ammonia  for  every  atom  of  platinum. 
The  chloride  in  this  series  is  therefore  chlor-platindiainine  chloride, 

and  has  the  formula  :  — 

„,  p    <(NH3)2C1 
C1»Ptj(NH3)2Cl. 


PLATINUM  ;    USES.  495 

In  this  compound  there  exists  the  divalent  radicle  C12  Pt  (  NH3  )4  =  , 
and  the  similar  combination  derived  from  platinous  chloride  (having 
two  less  chlorine  atoms),  has  the  formula :  — 

p   J(NH3)2C1 
tt(NH3)2Cl. 

One  other  oxide  of  this  family  should  be  mentioned.  This  is  the 
monoxide,  MO,  which,  as  well  as  the  dioxide,  is  common  to  all  the 
members  of  the  platinum  group  ;  it  is  basic  in  its  character,  and  from 
it  are  derived  the  -ous  salts. 

Platinous  chloride  has  the  formula  Pt  C12 .  It  is  formed  by  reduc- 
tion of  a  solution  of  platinic  chloride  by  means  of  sulphur  dioxide ; 
like  platinum  chloride,  it  readily  forms  double  salts,  and  unites  with 
ammonia. 

The  commercial  uses  of  the  metals  of  the  platinum  group  are 
confined  chiefly  to  the  preparation  of  standard  weights  and  meas- 
ures, and  to  the  manufacture  of  chemical  apparatus.  Most  of  the 
platinum  which  occurs  in  crucibles  and  other  utensils  is  not  pure ;  it 
contains  as  much  as  2  per  cent  of  iridium,  and,  by  very  reason  of 
this  admixture,  is  more  valuable,  because  less  readily  attacked  by 
acids.  Platinum  vessels  are  extremely  useful  because  of  their  high 
melting  point,  because  they  are  not  attacked  by  oxygen,  and  are  not 
dissolved  by  the  ordinary  laboratory  reagents.  In  using  platinum 
ware,  care  must  be  taken  not  to  fuse  caustic  alkalies  therein,  nor 
to  heat  with  any  metal  or  easily  reducible  compound  containing  a 
metal,  because  platinum  readily  forms  alloys.  Furthermore,  platinum 
unites  with  silicon  or  phosphorus,  the  compounds  so  formed  being 
very  brittle.  Charcoal,  or  coal  which  contains  silicon  compounds  in 
its  ash,  should  never  be  ignited  in  a  crucible  of  platinum ;  and  the 
.heating  of  compounds  which  contain  phosphates  in  the  presence  of 
a  reducing  agent  (such  as  charcoal)  should  also  be  avoided.  Lumi- 
nous gas  flames  render  platinum  rough  and  brittle,  the  same  is  true 
of  the  reducing  flame  of  a  Bunsen  burner.  Unless  care  is  taken  to 
remove  this  roughness  each  time  after  using,*  it  will  ultimately 
penetrate  the  crucible,  and,  after  rendering  the  utensil  brittle,  will 
crack  it. 


By  rubbing  with  a  little  fine  sand  until  smooth. 


APPENDIX  OF  LABORATORY  NOTES. 


BEFORE  entering  upon  laboratory  work,  the  pupil  should  read  and  remem- 
ber the  following  cautions :  — 

Burns,  Stains,  and  Fire.  —  Yellow  phosphorus  should  never  be  handled 
excepting  with  a  pair  of  tongs  or  pincers.  When  exposed  to  the  air  in  a 
warm  room,  it  may  take  fire  spontaneously;  if  touched  by  the  hand,  it  will 
take  fire;  the  burns  so  produced  are  extremely  painful,  and  may  become 
dangerous  by  reason  of  phosphorus  poisoning.  Sodium  and  potassium  are 
kept  under  naphtha;  they  should  never  be  placed  in  water  when  the  pieces 
are  larger  than  beans,  and,  in  any  event,  a  very  small  piece  should  first  be 
tested.  Sodium  which  has  not  a  clear  and  bright  surface  when  cut,  should  be 
rejected ;  in  all  cases  the  outer  coating  of  oxide  should  be  cut  away  before 
placing  the  metal  in  water.  Burns  are  best  treated  by  covering  the  spot  with 
a  solution  of  cocaine  in  sweet  oil,  and  then  with  an  emulsion  of  lime-water, 
glycerine,  and  sweet  oil.  Nitric  acid  stains  the  skin  yellow;  when  concen- 
trated, it  will  cause  an  ulcer  to  form.  Bromine  stains  the  skin  brown;  unless 
instantly  removed,  it  will  cause  a  painful  ulcer.  Iodine  stains  the  skin  dark 
brown ;  nitrate  of  silver,  black.  Of  course,  every  precaution  should  be  taken 
to  keep  the  above  substances  from  touching  the  hands  or  face ;  but,  in  case  of 
accident,  dilute  sodium  carbonate  solution,  followed  by  washing  with  clean 
water,  will  be  best  to  apply  in  the  case  of  nitric  acid  and  bromine.  A  solution 
of  sodium  hyposulphite  followed  by  water  will  remove  iodine;  cyanide  of 
potassium  solution  will  remove  silver  stains.*  Concentrated  sulphuric  acid 
and  solutions  of  chromic  acid  will  attack  the  skin,  not  so  rapidly,  however,  as 
nitric  acid ;  in  case  of  an  accident  they  can  often  be  removed  by  washing  with 
water  or  sodium  carbonate  solution  before  serious  results  have  followed.  Hot 
sulphuric  acid  will  instantly  produce  the  most  painful  burns;  any  test-tube  in 
which  sulphuric  acid  or  anything  else  is  being  heated  should  be  held  by  a  test- 
tube  holder,  and  with  its  mouth  pointing  away  from  the  manipulator  or  from 
any  one  standing  near.  Ether  or  carbon  bisulphide  must  not  be  used  on  a 
desk  within  at  least  six  feet  of  a  burning  Bunsen  burner,  as  these  liquids  take 
fire  with  the  greatest  readiness.  Matches  should  be  kept  in  a  tin  box,  which 
is  never  to  be  placed  in  the  drawer  of  the  desk,  but  should  always  be  kept 
outside. 

*  The  teachers  should  not  keep  cyanide  of  potassium  in  a  place  where  it  is  accessible 
to  any  one  but  themselves;  it  should  be  handled  under  a  hood,  and  should  be  applied  in 
removing  silver  stains,  to  the  hands  only,  and  then  be  instantly  washed  off  after  its  work  is 
accomplished. 

497  "°*  THS 


UJUYIRSITYl 


498  APPENDIX   OF 

Inhalation  of  Fumes  and  Gases.  —  Chemical  experiments  which  will  develop 
poisonous  or  irritating  gases  should  always  be  performed  under  a  hood  with  a 
good  draught.*  Chlorine,  bromine,  phosphorus  pentachloride,  phosphorus  tri- 
chloride, attack  the  mucous  membrane  of  the  eyes,  throat,  and  nose ;  continued 
inhalation  will  give  rise  to  bronchial  inflammation;  chlorine  or  bromine  will 
also  cause  nausea  to  ensue.  If,  by  accident,  the  pupil  should  take  an  exces- 
sive quantity  of  chlorine  into  the  lungs,  the  quickest  remedy  is  probably  the 
inhalation  of  the  fumes  of  alcohol.  The  gaseous  oxides  of  nitrogen  are 
poisonous,  causing  violent  headache  and  nausea;  phosphine,  arsine,  stibine, 
are  very  poisonous ;  ammonia  is  quite  irritating.  Work  in  which  these  sub- 
stances are  generated  or  used  must  be  done  under  the  hood.t  Sulphuric  acid 
should  not  be  heated  to  above  150°  unless  the  apparatus  is  under  the  hood. 
(The  acid  will  break  down,  partly,  into  H2O  and  SO3;  the  vapors  of  SO3  are 
irritating  to  the  lungs.)  Liquids  containing  hydrochloric,  nitric,  or  hydro- 
fluoric acids,  etc.,  should  be  evaporated  under  the  hood.  Sulphuretted  hydro- 
gen is  poisonous  and  disagreeable ;  continued  inhalation  of  even  small  quantities 
will  cause  headache,  and  may  have  serious  results.  It  is,  therefore,  imperatively 
necessary,  unless  a  room  is  especially  provided  in  which  to  generate  this  gas, 
that  all  work  with  hydrogen  sulphide  should  be  performed  under  the  hood. 

Explosions.  —  The  majority  of  accidents  result  from  carelessness,  there- 
fore the  invariable  rule  by  which  the  student  should  govern  himself  in  the 
laboratory  is,  "never  be  careless,  for  carelessness  may  result  in  permanent 
disfigurement  or  loss  of  sight."  Hydrogen  and  oxygen,  hydrogen  and  air, 
hydrogen  and  chlorine,  gaseous  hydrocarbons  and  oxygen,  phosphine  and 
oxygen,  or  phosphine  and  air,  as  well  as  the  other  not  very  stable  hydrogen 
compounds  of  the  not-metals  when  mixed  with  oxygen  or  air,  will,  unless  one  or 
the  other  constituent  is  present  in  proportionally  small  quantity,  cause  violent 
explosions  when  they  are  ignited.  In  generating  the  gases,  extreme  care  must 
be  taken  not  to  bring  a  flame  near  the  exit  tube  of  the  apparatus  until  the 
pupil  is  sure  that  a  brisk  current  of  the  generated  gas  has  traversed  the  appa- 
ratus for  sufficient  length  of  time  to  expel  all  air.  Of  course,  no  definite  time 
rule  can  be  established,  because  this  will  vary  with  the  size  of  the  apparatus; 
but  when  using  the  ordinary  generating  flasks  of  from  300  to  500  c.c.  the  pupil 
should  wait  at  least  7  to  10  minutes.  Chlorate  of  potassium,  permanganate  of 
potassium,  and  similar  powerful  oxidizers,  must  not  be  rubbed  in  a  mortar  when 
in  contact  with  substances  which  are  readily  oxidized  (sugar,  starch,  sulphide 
of  antimony,  sulphur,  phosphorus  [yellow  or  red],  etc.). 

In  a  well  conducted  laboratory,  desks  and  apparatus  are  always  kept  as 
clean  as  possible,  and  reagent  bottles  returned  to  their  proper  places  as  soon 
as  the  occasion  requiring  their  use  is  over.  Bunsen  burners  can  be  cleaned  by 

*  So  urgent  is  this  rule  that  pupils  should  be  forbidden  even  to  heat  test-tubes  or  small 
evaporating  dishes  with  reagents  which  will  give  off  fumes  of  hydrochloric  acid,  nitric  acid, 
hydrogen  sulphide,  bromine,  chlorine,  nitric  oxide,  etc.,  unless  they  do  so  under  the  hood. 
A  good  hood  is  as  necessary  as  a  good  burner. 

t  In  the  case  of  ammonia  the  precaution  may  be  omitted,  if  only  small  quantities  of  the 
reagent  are  to  be  used. 


LABORATORY   NOTES.  499 

unscrewing  the  outer  tube  and  brushing  the  nipple  with  a  dry,  stiff  test-tube 
brush.* 


The  numbers  of  the  following  notes  correspond  to  the  reference 
numbers  in  the  text. 

1.  PREPARATION  OF  OXYGEN  BY  HEATING  MERCURIC  OXIDE.  —  The 
oxide  decomposes  at  a  low  red  heat.     A  little  should  be  placed  in  a  glass  tube 
300  m.m.  in  length,  closed  at  one  end  and  made  of  so-called  infusible  glass  ; 
outside  of  the  tube,  place  a  cylinder  of  copper  wire  gauze  to  prevent  cracking; 
a  triple  gas-burner  is  most  convenient  for  heating.     The  tube  is  connected 
with  the  trough  of  water,  over  which  the  gas  is  collected,  by  means  of  a  safety 
bottle ;  the  latter  consisting  of  an  empty  4  oz.  wride-mouthed  bottle,  fitted  with 
a  double  bored  rubber  stopper  and  connecting  glass  tubes,  the  ends  of  which 
must  not  extend  below  the  bottom  of  the  stopper.     When  such  a  safety  bottle  is 
present,  the  water  in  the  trough  cannot  be  forced  back  into  the  hot  tube,  if,  by 
any  accident,  the  flame  should  be  extinguished ;  for  the  cold  water  cannot  get 
beyond  the  safety  bottle  ;  if  it  were  to  strike  the  hot  tube,  an  accident  would 
be  sure  to  follow.     Such  a  safety  bottle  should  always  be  interposed  where  a 
pneumatic  trough  is  used  to  collect  gases  which  are  generated  in  an  apparatus 
which  is  to  be  heated  to  a  high  temperature.     (In  the  apparatus  depicted, 
Fig.  1  (Page  19),  the  place  of  a  triple  burner  is  supph'ed  by  a  "combustion 
furnace,"  which  is  a  long  oven  heated  by  a  number  of  flames.) 

2.  As   BLACK  OXIDE  OF  MANGANESE  is  sometimes  adulterated  with 
charcoal,  it  is  always  necessary  to  test  the  chemical  by  heating  a  very  little 
of  it  in  a  test-tube  before  using  a  larger  quantity;  if  no  explosion  results,  it  is 
safe  to  use. 

3.  PREPARATION  OF  OXYGEN  BY  HEATING  MANGANESE  DIOXIDE. — 
The  apparatus  to  be  used  is  identical  with  that  used  in  the  experiment  (Note 
1),  excepting  that  an  iron  tube,  made  of   ordinary  gas-pipe,  capped  and  18 
inches  in  length,  is  substituted  for  the  glass  one.     Fill  this  tube  one-fourth  full 
of  manganese  dioxide,  broken  to  the  size  of  a  pea.t     Place  the  tube  flat  upon 
the  desk,  and  pound  sharply  after  filling;  this  is  for  the  purpose  of  making  a 
canal  for  the  passage  of  the  gas  above  the  load  in  the  tube.  J     Heat  to  a  red 
heat  in  a  combustion  furnace  (Fig.  1),  and  use  a  safety  bottle  (Note  1). 

4.  THE  CHLORATE  OF  POTASSIUM  should  be  tested  in  the  same  way  as 
the  black  oxide  of  manganese  (Note  2). 

*  Any  instruction  in  glass  bending  or  blowing  which  is  necessary  should  be  given  by 
the  teacher  before  beginning  laboratory  work;  after  the  instruction,  practice  alone  will 
make  perfect.  The  pupil  should  buy  W.  A.  Shenstone,  Methods  of  Glass  Blowing,  Riving- 
ton's,  London,  1886. 

t  If  this  cannot  be  procured,  take  powdered  magnanese  dioxide. 

£  This  precaution  must  never  be  omitted  in  charging  tubes  with  solid  substances  which 
are  to  generate  gases  on  heating;  otherwise  the  apparatus  will  certainly  explode. 


500 


APPENDIX   OF 


5.  PREPARATION  OF  OXYGEN  BY  HEATING  CHLORATE  OF  POTASSIUM.  — 
Take  a  flask  of  200  c.c.  capacity,*  fitted  with  a  single  bored  rubber  stopper  and 
glass  delivery  tube  (Fig.  15)  and  heat  to  a  low  red  heat,  collecting  the  gas 
over  water  in  a  pneumatic  trough ;  it  is  best  to  insert  a  safety  bottle  between 
the  generating  flask  and  the  water  (Note  1).  As  the  flask  in  which  this  oper- 
ation has  been  performed  is  always  incapacitated  for  future  use,  and  as  the 
operation  is  not  used  for  the  practical  preparation  of  oxygen,  it  is  better  to 

substitute  the  following:  Heat 
chlorate  of  potassium  in  a  hard 
glass  test-tube  until  all  of  the 
oxygen  is  expelled  ;  prove  that 
oxygen  is  present  by  placing  a 
glowing  pine  chip  in  the  tube. 
The  most  approved  method  of 
preparing  oxygen  for  laboratory 
use  is  by  heating  a  mixture  of 
chlorate  of  potassium  and  man- 
ganese dioxide.  Mix,  in  a  mor- 
tar, 25  grams  of  potassium 
chlorate  and  5  grams  of  man- 
ganese dioxide;  place  the  mix- 
ture in  a  flask  like  the  one 
indicated  in  the  first  part  of 
this  note  and  shown  by  Fig.  15 ; 
heat  gently  with  a  Bunsen 
burner  until  the  gas  conies  off 
slowly  and  regularly;  collect 
all  of  the  gas  in  bottles  over  t 
water  in  the  pneumatic  trough, 


Fig.  15. 


and  save  for  future  use. 


6.  COMBUSTION  IN  OXYGEN.  —  Burn  the  substances  mentioned  in  the 
text,  cutting  phosphorus,  sulphur,  and  carbon  to  the  size  of  a  pea.     The  de- 
flagrating spoon  in  which  these  substances  are  burned  (Fig.  16)  should  have 
its  handle   thrust  through  a  piece  of  sheet-zinc  larger  than  the  mouth  of 
the  jar  containing  the  gas  and  pierced  with  a  small  hole  in  the  centre.     It 
is  well  to  perform  the  burning  of  phosphorus  in  a  large  globe,  care  being 
taken  to  have  the  deflagrating  spoon  sink  so  far  into  the  vessel  as  to  reach 
the  centre;  if  the  jar  is  too  small,  the  heat  is  apt  to  crack  it. 

7.  COMBUSTION  OF  A  STEEL  WATCH-SPRING.  —  The  watch-spring  should 

*  Round-bottomed  flasks  made  of  hard  Bohemian  glass  have  lately  been  brought  into 
the  market;  they  are  in  every  way  more  desirable  than  flat-bottomed  or  Erlenmeyer  flasks. 

t  32  oz.  wide-mouthed  (so-called  salt  mouth)  common  bottles  are  cheapest  and  best  for 
collecting  gases.  The  pupil  should  have  a  number  of  square  pieces  of  window  glass  larger 
than  the  mouth  of  the  bottle;  when  necessary  to  remove  the  latter,  filled  with  gas,  from  the 
trough,  cover  the  mouth  with  one  of  the  pieces  of  glass,  by  pressing  the  same  against 
the  mouth  of  the  bottle,  raise  the  bottle  up  and  invert  it,  still  covered  with  the  glass. 


LABORATORY  NOTES. 


501 


be  heated  in  the  Bunsen  burner  and  then  straightened ;  a  small  piece  of  cotton 
is  now  tied  at  one  end  and  dipped  into  molten  sulphur;  on  igniting  the  sul- 
phur, and  plunging  the  spring  into 
a  jar  of  oxygen,  the  heat  given  off 
by  the  burning  sulphur  will  cause 
the  iron  to  take  fire.  A  cheap  jar 
should  be  employed ;  because,  during 
the  combustion,  the  temperature  rises 
high  enough  to  melt  the  oxide  of 
iron,  the  small  particles  of  which 
fly  off  and  become  fused  into  the 
walls  of  the  vessel,  and  may  even 
break  it.  To  illustrate  the  kindling 
temperature,  place  a  few  drops  of 
dry  carbon  bisulphide  in  a  test-tube, 
warm  until  the  tube  becomes  filled 
with  the  vapors  of  the  liquid,  and 
then  place  the  end  of  a  glass  rod, 
which  has  previously  been  warmed 
in  a  Bunsen  burner,  within  the 
mouth  of  the  test-tube;  the  carbon 
bisulphide  will  take  fire,  thus  furnish- 
ing an  illustration  of  a  low  kindling 
temperature;  the  pupil  can  find  ex- 


amples of  high  kindling  temperature 
without  suggestion. 


Fig.  16. 


8.  PREPARATION  OF  HYDROGEN  BY  MEANS  OF  SODIUM  AND  WATER.  — 
The   method  of  preparation  is   described  on  page  28  of  the  text,  and  the 
arrangement  of  the  apparatus  is  made  clear  by  figure  4  (page  29).     If  no  wire 
spoon  is  at  hand,  the  piece  of  sodium,  which  should  be  cut  to  about  the  size 
of  a  bean,  may  be  wrapped  in  a  small  piece  of  copper  gauze,*  and  this  piece 
then  taken  up  by  a  pair  of  crucible  tongs  and  slid  under  the  mouth  of  the 
test-tube,  which  has  previously  been  filled  with  water  and  inverted  in  the 
trough.    Care  should  be  taken  to  test  the  sodium  to  be  used  in  any  of  these  ex- 
periments by  placing  a  small  piece  on  water  and  then  standing  aside ;  for,  un- 
less the  metal  is  clean,  there  is  great  danger  of  an  explosion.     Scraps  of  sodium 
which  have  been  kept  in  the  laboratory  for  some  time  should  never  be  used. 

9.  The  method  for  preparing  hydrogen  by  means  of  sodium  and  water  is 
expensive,  and  not  to  be  recommended  for  ordinary  purposes.     It  has  been 
used,  however,  where  perfectly  pure  hydrogen  is  required. 

*  Filter-paper  will  even  answer  the  purpose.  If  copper  gauze  is  used,  the  sodium  and 
gause  will  sink  to  the  bottom,  the  sodium  will  melt  after  a  time,  and,  escaping  through  the 
meshes  of  the  gauze,  will  rise;  care  must  be  taken  to  place  the  tube  so  as  to  catch  the  parti- 
cles of  metal  as  they  rise.  If  paper  is  used,  the  sodium  and  paper  will  rise  to  the  top  of  the 
tube,  and  the  generation  of  hydrogen  will  go  on  at  the  point. 


502 


APPENDIX   OF 


10.  THE  DECOMPOSITION  OF  WATER  BY  POTASSIUM.  —  The  piece  of 
potassium  to  be  used  should  be  cut  about  the  size  of  a  pea  (handle  neither 
potassium  nor  sodium  with  the  fingers!),  and  thrown  on  the  water  in  the 
pneumatic  trough.     The  decomposition  of  water  by  potassium  is  so  violent 
that  the  heat  generated  sets  fire  to  the  hydrogen  evolved,  the  latter  burning 
with  a  violet  flame.     Care  must  be  taken  to  stand  at  some  distance  from  the 
water  on  which  the  potassium  is  floating,  as  an  explosion  may  occur  by  means 
of  which  the  pieces  of  the  metal  will  be  thrown  about.    Such  pieces  may  cause 
painful  burns,  and,  if  in  contact  with  the  eyes,  possible  loss  of  sight. 

11.  GRANULATED  ZINC. — Prepared  by  fusing  zinc  in  a  stoneware  cru- 
cible, and  pouring  the  melted  metal  in  a  thin  stream  into  cold  water,  care 
being  taken  to  stir  the  water  vigorously  during  the  operation. 

12.  To  PREPARE  HYDROGEN  BY  MEANS  OF  ZINC  AND  DILUTE  SUL- 
PHURIC ACID.*  —  Place  5  grams  of  zinc  in  a  gas-generating  flask  fitted  with 

a  double  bored  rubber  stop- 
per and  delivery  and  safety 
tube  (Fig.  17)  ;t  now  pour 
dilute  sulphuric  acid  through 
the  safety  tube  on  to  the 
zinc,  adding  acid  from  time 
to  time  as  occasion  requires. 
Dilute  sulphuric  acid  is  pre- 
pared by  adding  one  part  of 
commercial  acid  to  six  parts 
of  water;  in  diluting  sul- 
phuric acid,  pour  the  acid 
gradually  into  the  water,  but 
do  not  pour  the  water  into 
the  acid ;  cool  the  acid  before 
using,  by  placing  the  flask 
under  the  hydrant. 

13.  IN  ORDER  TO  PU- 
RIFY THE  HYDROGEN  pre- 
pared as  in  note  12,  it  should 
be  passed  through  a  train  of 
wash-bottles  (Fig.  18).  In 
the  first  is  a  solution  of  po- 
Fig,  17.  tassium  permanganate  to  re- 

move gases    which   can    be 
oxidized;   viz.,  hydrocarbons,  hydrogen  sulphide,  hydrogen  arsenide,  etc.;  in 


*  Be  sure  to  use  zinc  which  is  free  from  arsenic,  or  which,  at  best,  contains  only  a  trace 
of  that  element. 

t  A  thick  walled  flask  of  300  c.c.  capacity  can  be  used  for  this  and  subsequent  opera- 
tions for  generating  a  gas,  provided  there  is  no  necessity  of  heating  the  flask. 


LABORATORY  NOTES. 


503 


the  second,  potassium  hydroxide,  to  remove  acid  vapors;  and  in  the  third, 
concentrated  sulphuric  acid,  to 
remove  moisture.  Such  a  train 
of  bottles  can  be  applied  wher- 
ever it  is  necessary  to  purify  a 
gas,  provided,  always,  the  sub- 
stances used  in  them  are  varied 
according  to  the  nature  of  the 
latter,  so  that  no  decomposition 
of  the  gas  can  result.  A  judi- 
cious choice  of  these  washing 
agents  comes  with  larger  expe- 
rience. 

14.    WHEN  LARGER  QUAN- 
TITIES OF  GAS  ARE  REQUIRED,  Fig.  18. 
it  is  better  to  use  a  form  of  gas 

generator  depicted  in  Fig.  19.  A  vessel,  6  d,  constricted  in  the  middle,  is 
fitted  with  a  glass  stopcock  delivery  tube,  e,  and  a  wide  globe  funnel,  with 
long  stem,  is  placed  in  the  upper  opening.  The  zinc 
is  placed  in  6,  and  the  acid  added  from  above  until 

the  apparatus  is  filled  to  about 

the  middle  of  the  funnel;  on 

opening  the  stopcock  the  acid 

ascends  to  the  metal ;  on  clos- 
ing, the  generated  hydrogen 

once    more    expels    the    acid 

from  the  central  globe.      In 

this   way  the  metals   can  be 

kept  indefinitely  out  of  con- 
tact with  the  acid,  and  need 

only  be  acted  on  by  it  when 

the  stopcock  e  is  opened.  This 

form   of    gas-generating    ap- 
paratus is  known  as   Kipp's 

gas  generator.     Several  other 

forms  have  been  devised,  but 

this  one  seems  to  be  the  most 

satisfactory  ;   it  is  useful  for 

generating  any  gas  which  does 


Fig.  19. 


Fig.  20. 


not  require  heat  in  its  manufacture.  Never  pour  hot 
acid  into  a  Kipp  generator,  nor  ever  pick  it  up  other- 
wise than  by  grasping  it  with  both  hands  around  the 
central  globe. 

Gases  which  are  desired  for  future  use  are  stored  in  a  gasometer  (Fig.  20). 
A  lower  metal  or  glass  tank,  a,  holding  about  forty  litres,  is  connected  with  an 
upper  one,  6,  by  means  of  two  tubes,  one  of  which  reaches  to  the  bottom  of  a. 


504 


APPENDIX   OF 


The  gasometer  is  filled  with  water,  all  stopcocks  closed,  the  cap  covering  the 
bottom  opening  is  removed,  and  the  gas  to  be  stored  is  run  in  through  this; 
when  all  the  water  has  been  replaced,  the  cap  is  screwed  on.  When  the  gas  is 
to  be  used,  the  upper  tank  is  filled  with  water,  the  stopcock  on  the  tube  lead- 
ing to  the  bottom  is  opened,  and  the  gas  allowed  to  escape  through  the 
upper  side  opening,  as  required. 

15.  DIFFUSION  OF  HYDROGEN  THROUGH  A 
POROUS    SUBSTANCE.  —  Construct   an   apparatus 
such  as  is  shown  in  Fig.  21.    This  consists  of  a  clean, 
dry,  porous  cup,*  which  is  fastened  at  the  end  of 
the  tube  c  by  means  of  a  funnel ;  this  funnel  has 
exactly  the  diameter  of  the  mouth  of  the  cup,  and 
the  two  are  fastened  together,  air-tight,  by  means 
of  rubber  cement  placed  around  the  rims.      The 
tube  c  is  connected  with  one  opening  of  a  double- 
necked  flask  t  by  means  of  a  single  bored  rubber 
stopper,  while  a  glass  tube,  drawn  to  a  point  and 
reaching  to  the  bottom  of  the  flask,  is  connected 
with  the  other  opening,  also  by  means  of  a  rubber 
stopper.     When  all  the  connections  are  air-tight, 
place  the  porous  cup  in  an  atmosphere  of  hydro- 
gen, by  inverting  over  the  cup  a  glass  bell,  open  at 
one  end  and  connected  with  a   hydrogen   gener- 
ator; the  hydrogen  will  rush  through  the  porous 
cup  much  more  rapidly  than  the  air  can  escape, 
for  the  specific  gravity  of  air  is  14.4  times  that  of 
hydrogen.     As  a  consequence,  the  air  in  the  cup 
and  in  c  is  forced  down  as  if   by  a  piston;  this 
causes  a  pressure  on  the  water  J  which  has  previ- 
ously been  placed  in  the  flask,  and  a  fountain  en- 
sues.    For  an  apparatus  to  illustrate  effusion  quan- 
titatively, see  Freer;  Zeitschrift  fur  Physikalische 
Chemie;   IX.  669. 

16.  OCCLUSION  OF  HYDROGEN.  —  Palladium 
is  very  expensive ;  one  gram  will  suffice  to  show  the 
occlusion  of  hydrogen.     Attach  the  palladium  to  a 
platinum  wire,  heat  red-hot  in  a  Bunsen  burner, 

allow  to  cool  slightly,  and  then  place  in  a  current  of  hydrogen  passing  from 
a  generator  ;§  the  occluded  hydrogen  will  be  oxidized  by  the  air,  the  palladium 
will  once  more  begin  to  glow,  and  will  finally  heat  the  hydrogen  to  its  kindling 
temperature.  So-called  platinized  asbestos  is  cheaper  and  just  as  available  for 
this  experiment.  Prepare  platinized  asbestos  by  heating  asbestos  which  has 
been  dipped  in  a  solution  of  platinum  chloride  and  then  into  ammonium 

*  A  small  porous  cup  from  a  Bunsen  battery. 

t  So-called  Woulff's  bottle. 

t  For  lecture-room  demonstrations,  color  the  water  with  blue  litmus. 

§  Take  care  to  have  all  of  the  air  expelled  from  the  generator! 


Fig.  21. 


LABORATORY   NOTES. 


505 


chloride.  A  little  of  this  may  be  fastened,  by  means  of  fine  platinum  wire,  within 
a  small  loop  made  of  iron  wire,  to  which  a  handle  about  four  inches  in  length 
is  attached.  Be  sure  to  perform  the  simple  experiment  given  on  the  bottom  of 
page  34,  and  the  last  one  on  the  first  paragraph  of  the  same  page. 

17.  TO  PROVE  THAT  HYDROGEN  FORMS  WATER  WHEN  BURNED  IN  AlR. 

—  Dry  the  gas  passing  from  the  generator  by  means  of  the  train  (Note  13),  or 
by  passing  it  through  a  glass  U-tube  filled  with  coarse  fragments  of  brick  soaked 
with  concentrated  sulphuric  acid,  or  through  a  similar  tube  filled  with  frag- 
ments* of  granulated  calcium  chloride  (Fig.  22). 

18.  Care  should  be  taken  to  expel  all  of  the  air  from  the  apparatus  before 
lighting.     To  find  out  if  this  has  been  accomplished,  test  the  gas  by  collecting  a 
test-tube  full  over  water  in  a  pneumatic  trough,  and  lighting  it;  if  it  burns 
quietly,  the  apparatus  is  safe.    The  burner  can  be  made  by  drawing  a  glass  tube 
nearly  to  a  point,  inserting  a  small  cylinder  of  rolled  platinum  foil  in  the 
small  end  so  produced,  and  then  fusing  the  glass  around  the  platinum  ;   after 

the  hydrogen  burner  has  • 

been  lighted,  place  a  large 

cold  beaker  over  the  flame, 
and  water  will  collect  there- 
in. 


19.  AN    EXPLOSIVE 
MIXTURE  OF  HYDROGEN 
AND  OXYGEN,  —  This  can 

be  prepared  with  safety  by  Fig  22. 

using  a  small  bottle,  with- 
out a  bottom  and  with  a  narrow  neck;  to  the  latter,  by  means  of  a  single  bored 
rubber  stopper  and  short  glass  tube,  a  long  rubber  tube  is  fitted.  This  tube 
can  be  closed  by  means  of  a  pinchcock.  The  bottle,  after  the  rubber  tube  is 
closed  by  the  pinchcock,  is  filled  with  water  in  the  pneumatic  trough,  and  then 
hydrogen  is  run  in  until  |  of  the  water  has  been  expelled ;  the  remaining  J  is 
then  similarly  replaced  by  oxygen.  The  mixture  of  gases  can  be  expelled  by 
lowering  the  bottle  in  the  pneumatic  trough  and  opening  the  pinchcock.  By 
placing  the  end  of  the  rubber  delivery  tube  under  some  soap-suds  in  a  small 
tin  dish,  a  few  soap-bubbles  filled  with  oxy-hydrogen  can  be  produced,  and 
then  exploded  by  touching  with  a  lighted  taper.  Care  should  be  taken  to 
remove  this  dish  to  a  safe  distance  from  the  bottle  before  exploding,  and  what- 
ever oxy-hydrogen  is  left  should  be  allowed  to  escape  as  soon  as  the  experi- 
ment requiring  its  use  has  been  completed.  If  possible,  some  experiments 
with  the  oxy-hydrogen  blow-pipe  (page  35)  should  be  performed. 

20.  THE  VOLUMETRIC  COMPOSITION  OF  WATER.  —  Use  a  eudiometer 
holding  50  c.c.     This  instrument  is  a  heavy  glass  tube  (Fig.  23),  which  is  closed 
at  one  end  and  graduated  in  millimetres  (sometimes  in  cubic  centimetres),  and 
has  two  platinum  wires  inserted  near  the  closed  tip  in  such  a  manner  that  an 


The  size  of  a  pea. 


506 


APPENDIX   OF 


Fij.  23. 


electric  spark  can  pass  from  one  to  the  other.*  This  tube  is  filled  with  mer- 
cury and  inverted  over  a  mercury  trough.  In  order  to  prove  the  volumetric 
composition  of  water,  slant  the  tube  to  one  side,  and  admit  about  lOc.c.  of 

hydrogen,  then  add  7  c.c.  of  oxygen  ; 
clamp  the  tube  tightly  with  its  open 
end  forced  against  a  leather  washer 
in  the  bottom  of  the  mercury  trough, 
taking  care  to  note  the  exact  volume 
of  hydrogen  admitted,  the  volume  of 
oxygen,  the  temperature,  barometer, 
and  height  of  the  mercury  in  the 
tube.  After  all  of  these  prepara- 
tions are  completed,  a  spark  from 
an  induction  coil  is  passed  through 
the  gas.  After  the  explosion  is 
completed  t  and  the  apparatus  has 
cooled,  note  the  volume  of  gas  re- 
maining, the  temperature,  and  the 
barometric  pressure,  as  well  as  the 
height  of  the  column  of  mercury  in 
the  tube.  Reduce  the  gas  before  and 
after  explosion  to  standard  condi- 
tions by  means  of  the  formula  on 
page  173,  remembering  that  the  height  of  the  column  of  mercury  in  the  tube, 
before  and  after  explosion,  must  be  deducted  from  the  barometric  pressure.  If 
the  pupil  used  10 c.c.  of  hydrogen  and  7  c.c.  of  oxygen,  then  the  10 c.c.  of 
hydrogen  will  have  united  with  5  c.c.  of  oxygen  to  form  water,  and  2  c.c.  of 
oxygen  will  remain.  The  decomposition  of  water  by  the  electric  current  is  made 
clear  by  Fig.  3.  The  two  electrodes  are  pieces  of  platinum  foil  ;  the  water  to 
be  decomposed  is  acidulated  with  sulphuric  acid. 

21.  THE  COMPOSITIOX  OF  WATER  BY  WEIGHT.  —  Dumas  passed  hydro- 
gen through  a  series  of  U-tubes  (Fig.  22)  filled  as  follows,  counting  from  his 
hydrogen  generator  (or  gasometer):  — 

U-tube  No.  1 ;  glass  fragments  moistened  with  lead  nitrate  solution,  to  remove 
hydrogen  sulphide. 

U-tube  No.  2;  glass  fragments  moistened  with  silver  sulphate  solution,  to 
remove  hydrogen  arsenide. 

U-tube  No.  3;  pumice-stone  moistened  with  caustic  potash  solution,  to  remove 
carbon  dioxide,  and  other  acids,  such  as  sulphurous  acid,  hydrochloric  acid,  etc. 

U-tubes  No.  4,  5;  pieces  of  solid  caustic  potash,  to  remove  acids. 

U-tubes  No.  6,  7,  8;  phosphoric  anhydride,  to  remove  moisture. 

The   last  tube  he  weighed  before   and   after  the   operation  ;   if   it  had 

*  Not  infrequently  the  eudiometers  come  to  the  laboratory  with  the  wires  so  close 
-together  that  the  spark  will  not  be  large  enough  to  ignite  the  gases;  if  such  is  the  case,  force 
the  ends  apart  with  a  long  glass  rod. 

~t  Wrap  a  towel  around  the  eudiometer  tube  before  exploding. 


LABORATORY   NOTES. 


507 


changed  in  weight  at  all,  the  experiment  was  rejected,  as  the  hydrogen  was 

not  pure.    The  hydrogen 

so  purified  was  passed 
through  a  tube  of  hard 
glass,  containing  copper 
oxide ;  this  tube  was  evac- 
uated by  an  air-pump 
before  beginning  the  op- 
eration, and  weighed 


Fig.  22. 


carefully  ;  it  was  then 
connected  at  one  end 
with  the  apparatus  delivering  hydrogen  which  was  purified  as  above  ;  at  the 
other  end,  with  two  U-tubes  filled  with  phosphorus  pentoxide  ;  the  tubes  were 
carefully  weighed  before  the  operation.  The  hydrogen  was  now  passed  over 
the  copper  oxide  in  the  tube,  which  was  heated  in  a  combustion  furnace 
(Fig.  1)  until  the  oxide  was  reduced  to  metallic  copper.  After  the  reduction 
was  complete,  and  the  tube  had  completely  cooled,  it  was  once  more  evacuated 
and  weighed  ;  the  loss  in  weight  was  equal  to  the  weight  of  oxygen  used  to 
form  water.  The  two  phosphorus  pentoxide  tubes,  which  were  placed  after  the 
copper  oxide,  were  also  weighed;  the  gain  in  weight  was  equal  to  the  amount 
of  water  which  had  been  formed.  The  pupil  should  perform  this  operation 
with  a  simpler  apparatus,  which  can  consist  of  the  train  of  wash-bottles 
(Note  13,  Fig.  18)  filled  (counting  from  the  hydrogen  generator),  No.  1  with 
potassium  permanganate  solution,  No.  2  with  lead  nitrate  solution,  No.  3 
with  caustic  potash  solution,*  and  lastly  a  U-tube  filled  with  pieces  of  brick 
soaked  in  sulphuric  acid.  The  copper  oxide  tube  should  be  of  hard  glass,  with 
a  bulb  blown  on  the  middle,  and  one  end  drawn  to  a  narrow  opening.  To  this 
end,  after  filling  the  bulb  with  granulated  copper  oxide  (the  copper  oxide 
should  previously  be  dried  at  150°  and  introduced  into  the  tube  while  still 
warm),  and  weighing  the  tube  on  the  balance  (weigh  tOy^  of  a  milligram), 
attach  a  U-tube  filled  with  pieces  of  brick  soaked  in  sulphuric  acid  (weigh 
this  tube  also  to  Ta^  of  a  milligram).  Now  pass  the  hydrogen  through  the 
apparatus,  while  heating  the  copper  oxide  to  a  low  red  heat.  After  reduction, 
is  complete,  continue  to  pass  hydrogen  for  some  time,  so  that  every  trace  of 
water  is  carried  over  into  the  last  U-tube  ;  allow  to  cool,  weigh,  and  calculate 
the  results  as  given  on  page  38.  Unless  the  ratio  is  within  of  .05  a  unit  of  the 
proportion  1  :  8,  the  experiment  should  be  repeated.  The  connections  be- 
tween the  various  parts  of  the  apparatus  are  made  by  means  of  rubber 
tubing  ;  be  sure  to  have  this  small  enough  to  fit  air-tight  over  the  ends  of  the 
tubes,  and  make  the  connections  as  short  as  possible. 

To  PROVE  THE  PRESENCE  OF  WATER  OF  CRYSTALLIZATION.  —  Heat  a 
crystal  of  copper  sulphate  in  a  test-tube,  and  see  if  water  passes  off,  whether 
at  a  high  or  low  temperature  ;  next  dissolve  the  anhydrous  salt  in  water, 

*  The  wash-bottles  cannot  be  filled  with  very  much  of  the  required  solutions,  otherwise 
the  pressure  in  the  hydrogen  generator  will  not  be  sufficient  to  overcome  the  hydrostatic 
pressure  in  the  wash-bottles ;  the  gas  would,  in  consequence,  not  flow  through  the  bottles. 


508  APPENDIX  OF 

evaporate  in  a  porcelain  dish  on  a  water-bath  *  until  crystallization  begins, 
set  aside  to  cool,  and  observe  the  form  of  the  crystals.  Note  the  change  in 
temperature  caused  by  dissolving  sodium  nitrate  in  water,  and  by  dissolving 
fused  calcium  chloride  in  water;  take  a  small  piece  of  quick-lime  and  place  it 
in  a  porcelain  dish,  then  pour  on  a  little  water  and  note  the  change  in 
temperature.  (See  pages  43,  44.)  An  example  of  an  efflorescent  salt  can  be 
obtained  by  procuring  a  few  "soda  crystals"  (Na2CO3 -f  10  H2O)  ;  a  deli- 
quescent substance  is  fused  calcium  chloride. 

22.  PREPARATION   OF  OZONE.  — The  apparatus,  Fig.  24,  is  the  one 
generally  employed  for  obtaining  considerable  quantities  of  ozone.     An  outer 

coating,  a,  is 
of  tinfoil,  sur- 
rounding a 
glass  tube; 
within  an 
inner  glass 

L — '      tube  (c)  is  a 

Fig.  24  piece  of  thin 

copper  foil ; 

the  two  glass  tubes  are  fused  together  at  &,  so  that  c  is  within  a  and  reaches 
nearly  to  the  end  of  the  latter  farthest  from  6.  Dry  oxygen  is  admitted  to 
the  space  between  a  and  c  by  the  tube  at  6;  it  passes  out  at  the  farther  end;  a 
is  connected  by  means  of  a  metal  strip  with  one  pole  of  an  induction  coil ;  c, 
by  means  of  a  similar  metal  strip  placed  at  6,  with  the  other  pole.  When  a 
current  passes,  a  silent  discharge  takes  place  between  the  tin  foil  surround- 
ing a  and  the  copper  contained  in  c.  This  discharge  must  necessarily  traverse 
the  oxygen  which  is  passing  in  at  6,  and  by  this  means  ozone  is  generated. 
The  preparation  of  ozone  by  the  student  can  be  accomplished  by  placing  a 
few  pieces  of  phosphorus  in  a  good  sized  bottle  with  a  wide  neck,  covering 
them  partly  with  a  very  dilute  aqueous  solution  of  potassium  dichromate  and 
.-sulphuric  acid,  warming  the  whole  slightly  (to  about  24°  —  about  the  tempera- 
ture of  a  hot  room  in  summer),  and  shaking  gently  from  time  to  time.  The 
bottle  will  soon  contain  ozone  ;  the  latter  can  be  detected  by  the  odor  and  by 
hanging  a  strip  of  paper,  soaked  with  a  mixture  of  starch  paste  and  a  little 
potassium  iodide,  within  the  neck  of  the  bottle  (page  50). 

23.  The  blue  color  of  ozone  may  be  seen  by  filling  a  glass  tube,  1  metre  in 
length,  with  ozonized  oxygen  and  then  looking  through  this  against  a  white 
background. 

24.  PREPARATION  OF  A  SOLUTION  OF  HYDROGEN  DIOXIDE.  —Add  finely 
powdered  barium  dioxide,  gradually,  to  dilute  sulphuric  acid  until  a  sufficient 
quantity  has  been  added  to  convert  all  of  the  sulphuric  acid  into  barium  sul- 

*  A  copper  basin,  hemispherical  in  shape,  with  the  top  formed  of  concentric,  movable, 
copper  rings.  This  vessel  is  filled  with  water  and  heated  with  a  Bunsen  burner,  substances 
to  be  evaporated  are  placed  in  porcelain  evaporating  dishes  upon  it.  Remember  that  a 
flame  merely  large  enough  to  cause  the  water  to  boil  is  just  as  good  for  heating  purposes  as 
one  which  will  cause  violent  ebullition,  and  which  will  necessitate  frequent  refilling  of  the 
water  bath. 


LABORATORY   NOTES. 


509 


phate,*  taking  care  to  keep  the  liquid  quite  cool  during  the  operation,  as 
warmth  destroys  hydrogen  hyperoxide  ;  the  barium  sulphate  is  allowed  to 
settle;  the  nearly  clear,  supernatent  liquid  is  poured  off  and  finally  filtered. 
To  a  portion  of  the  colorless  liquid  so  formed,  add  a  mixture  of  starch  paste 
and  a  little  iodide  of  potassium  solution,  and  observe  the  blue  color  of  iodine- 
starch  (page  50);  to  another  portion  add  a  dilute  solution  of  potassium  di- 
chromate  and  a  little  sulphuric  acid;  an  intensely  blue  color  (ascribed  to 
the  formation  of  perchromic  acid)  will  be  seen;  shake  the  blue  solution  in  the 
test-tube  with  a  little  ether,  t  and  the  blue  substance  will  be  dissolved  by 
the  latter. 

25.  The  solution  of  hydrogen  peroxide  is  concentrated  by  placing  it  in 
an  open  dish  under  the  bell  of  an  air-pump  which  also  has  under  it  an  open 
dish  containing  concentrated  sulphuric  acid;  the  air  is  then  exhausted,  and  the 
solution  allowed  to  evaporate.     This  is  a  common  method  for  concentrating 
liquids  which  are  easily  decomposed. 

26.  ELECTROLYSIS  OF  A  SOLUTION  OF   HYDROCHLORIC  ACID.  —  The 
apparatus  to  be  used  is  shown  by  Fig,  25  ;  it  explains  itself,  with  the  exception 
that  the   two  elec- 
trodes   within    the 

U-shaped  tube  must 
be  made  of  gas  car- 
bon and  not  of  plat- 
inum. Such  an 
apparatus  may  be 
especially  ordered 
for  this  experiment, 
or  the  apparatus' 
given  by  Fig.  3  can 
be  used,  the  two 
platinum  electrodes 
being  removed,  and 
in  their  places  two 
small  pieces  of  gas 
carbon,  the  con- 
nections made  by 
platinum  wire,  can 
be  inserted.  In- 
stead of  an  ordi- 
nary solution  of 
hydrochloric  acid, 
a  solution  of  com-  Fig.  25. 

mon  salt,  saturated 

*  This  point  can  be  ascertained  by  allowing  the  solution  to  settle,  taking  out  a  drop  of 
the  clear  liquid,  and  adding  this  to  a  drop  of  barium  chloride  solution  on  a  watch-glass;  if  no 
insoluble  barium  sulphate  is  formed,  the  sulphuric  acid  has  all  been  converted  to  barium 
sulphate. 

t  By  pouring  a  little  ether  over  the  liquid,  placing  the  thumb  over  the  mouth  of  the  test- 
tube,  and  shaking. 


510 


APPENDIX   OF 


at  ordinary  temperatures,  to  which  ^  of  its  volume  of  concentrated  hydro- 
chloric acid  has  been  added,  is  used.  When  the  electric  current  is  turned  on, 
the  chlorine  will  at  first  be  absorbed  by  the  liquid  which  fills  the  apparatus; 
after  the  latter  is  entirely  saturated  with  the  gas,  then  equal  volumes  of 
hydrogen  and  chlorine  will  be  produced  simultaneously.  Care  must  be  taken 
not  to  open  the  stopcocks  too  frequently  while  the  operation  is  going  on, 
otherwise  the  liquids  in  the  two  arms  of  the  U-shaped  tube  become  mixed,  and 
the  action  irregular.  Too  much  liquid  must  not  be  placed  in  the  apparatus,  as, 
otherwise,  the  increase  in  pressure,  caused  by  the  liberation  of  the  gases  and 
consequent  rising  of  liquid  in  the  rear  tube,  will  cause  a  greater  absorption  of 
chlorine. 

27.  THE  PREPARATION  OF  CHLORINE.  —  The  apparatus  best  adapted 
for  the  preparation  of  all  gases  which  are  either  noxious  or  poisonous  is 
shown  by  Fig.  26.  When  the  gas  is  desired  for  use,  the  stopcock  B  is  opened 

and  A  is  closed  ;  the  gas 
can  then  pass  from  the 
generating-flask  into  the 
vessel  or  apparatus  where 
£  it  is  to  be  used.  When  the 

gas  is  not  required,  the 
stopcock  B  is  closed  and 
A  is  opened ;  then  it  passes 
from  the  generating-flask 
into  a  vessel  filled  with 
some  liquid  which  will 
completely  absorb  it.  For 
the  preparation  of  chlorine 
use  a  generating  -  flask 
(round  bottomed  preferred) 
of  about  500 c.c.  capacity; 
place  in  this  50  grams  of 
manganese  dioxide  (pow- 
dered); unite  all  parts  of 
the  apparatus,  and  then 
pour  in  a  sufficient  quan- 
tity of  hydrochloric  acid 
through  the  safety  tube  to 
cover  the  manganese  dioxide ;  then  warm  gently  with  a  Bunsen  burner ;  or, 
mix  in  a  mortar  50  grams  of  manganese  dioxide,  50  grams  of  sodium  chlo- 
ride, place  in  the  generating-flask,  connect  all  parts  of  the  apparatus,  and 
then  add  150  c.c.  of  sulphuric  acid  (two  parts  sulphuric  acid  to  one  of  water), 
and  warm  gently.  The  pupil  should  prepare  chlorine  by  both  of  these 
methods.  Collect  the  chlorine  in  dry  bottles  by  displacement  of  air,  as  is 
shown  in  the  cut,  and,  when  full,  cover  the  bottles  with  glass  plates,  and  place 
aside  for  future  use  ;  afterwards  pass  chlorine  into  300  c.c.  of  water  until 
the  liquid  is  saturated  with  the  gas.  Form  chlorine  hydrate  crystals,  as  indi- 


Fig.  26. 


LABORATORY   NOTES. 


511 


cated  on  page  62  (paragraph  1).      Into  the  jars  filled  with  chlorine  intro- 
duce :  — 

a.  A  little  powdered  antimony. 

b.  A  few  pieces  of  heated  copper  foil. 

c.  A  piece  of  moist  litmus  paper. 

d.  A  piece  of  filter  paper,  moistened  with  turpentine. 

Try  the  bleaching  power  of  your  chlorine  water  on  a  piece  of  colored  calico. 
Experiments  with  chlorine  must  be  conducted  under  a  hood ! 

28.  BLEACHING  BY  MEANS  OF  CHLORINE.  —  This  can  best  be  shown  by 
passing  dry  chlorine  *  through  a  flask  of  1  litre  capacity  which  has  been  fitted 
with  a  triple  bored  rubber  stopper;  connect  one  of  the  glass  tubes  introduced 
through  this  with  the  chlorine  generator;  connect  the  second  with  a  small 
flask  containing  water,  so  that  the  latter  can  be  boiled  when  desired,  so  as  to 
force  steam  into  the  apparatus ;  place  a  glass  tube  in  the  third  hole  of  the 
triple   bored   stopper   (this  last  is  for  the  escape  of  the  superfluous  gases); 
it  can,  if  desired,  be  so  bent  as  to  open  into  a  small  jar  containing  caustic 
soda  solution.     Introduce  a  piece  of  colored  calico  into  the  1  litre  flask,  and 
admit  dry  chlorine:  no  bleaching  action  will 

be  observed.  Now  heat  the  water  in  the  small 
flask  and  force  in  steam:  the  calico  will  then 
be  bleached. 

29.  PREPARATION    OF    HYDROCHLORIC 
ACID.  —  The   apparatus   to    be    used    is    the 
same  as  that  for  the  preparation  of  chlorine 
(Fig.  26).      The  generator  is  charged  with  20 
grams  of  sodium  chloride,  and  then  30  grams 
of  sulphuric  acid  (two  parts  of  acid  to  one  of 
water)  are  added   through   the   safety  tube; 
heat  very  gently,  and  collect  two  jars  of  the 
gas  as  was  done  with  chlorine,  and  pass  the 
remainder  of  the  gas  into  a  beaker  contain- 
ing water. 

Hydrochloric  acid  is  produced  by  burning 
chlorine  in  an  atmosphere  of  hydrogen.  —  In 
a  small  flask  (A,  Fig.  27)  of  150 c.c.  capacity, 
place  10  grams  of  powdered  manganese  diox- 
ide. Fit  the  flask  with  a  single  bored  rubber 
stopper  into  which  is  inserted  a  tube,  widened 
at  the  centre,  and  with  its  widened  part  filled 
with  pieces  of  granulated  calcium  chloride 
about  the  size  of  a  pea;  after  pouring  con- 
centrated hydrochloric  acid  on  the  manganese 


Fig,  27. 


dioxide,  put  this  stopper  into  the  flask,  allow  the  chlorine  to  escape  for  a  time, 
and  then  invert  a  jar,  #,  of  hydrogen  over  the  escaping  chlorine,  taking  care  to 

*  Dry  the  chlorine  by  passing  it  through  a  U-tube  containing  pieces  of  brick  saturated 
with  cone,  sulphuric  acid  (Note  21); 


512 


APPENDIX   OF 


light  the  hydrogen  just  before  it  reaches  the  chlorine  jet;  the  chlorine  will 
then  burn  in  the  hydrogen  which  fills  the  jar.  The  reverse,  burning  hydrogen 
in  chlorine,  can  be  performed  by  filling  the  flask  A  with  zinc  and  dilute  sul- 
phuric acid,  and  lengthening  and  bending  the  delivery  tube,  so  that  it  may 
form  a  burner  which  can  extend  down  into  a  jar  of  chlorine ;  after  hydrogen 
has  expelled  all  of  the  air  from  the  apparatus,  light  the  jet,  and  lower  the 
flame  into  the  chlorine  jar;  it  will  continue  to  burn.  In  both  experiments 
fumes  of  hydrochloric  acid  will  be  observed. 

30.  EXPERIMENTS  WITH  HYDROCHLORIC  ACID.  — Place  a  strip  of  moist- 
ened blue  litmus  paper  in  one  of  the  jars  of  hydrochloric  acid  gas ;  place  a 
lighted  taper  in  the  other,  and  see  if  the  gas  supports  combustion.    To  a  little  of 
your  solution  in  a  test-tube  add  some  iron  filings ;  some  pieces  of  zinc ;  dilute 
the  solution  and  taste  it;  try  its  effects  on  blue  litmus  solution.    To  some  dilute 
hydrochloric  acid,  to  which  you  have  added  a  few  drops  of  litmus  solution, 
add  a  solution  of  sodium  hydroxide,  drop  by  drop  from  a  burette,*  until  the 
red  color  of  the  litmus  just  turns  into  blue ;  one  drop  of  acid  will  then  turn  the 
blue  litmus  red,  a  drop  of  sodium  hydroxide  will  turn  the  color  back  to  blue; 
the  solution  is  then  neutral.      If  it  is  evaporated,  nothing  but  sodium  chloride 
will  remain ;  that  sodium  chloride  is  formed  can  be  proved  by  adding  sulphuric 
acid  to  the  crystallized  remainder.     (Pages  76  and  77.) 

31.  The  electrolysis  of  hydrochloric  acid  is  described  in  Note  20. 

The  extreme  solubility  of  hydrochloric  acid  in  water  may  be  shown  by  an 
apparatus  such  as  is  depicted  by  Fig.  28.  The  upper  flask  is  well  filled  with 
dry  hydrochloric  acid  gas,  and  is  stoppered  with  a  single  bored  rubber  stopper ; 

through  the  stopper  there  runs  a 
glass  tube,  drawn  to  a  point,  and 
sealed  without  the  flask,  and  nar- 
rowed within.  The  large  beaker 
is  filled  with  blue  litmus  solution. 
When  all  is  ready,  the  sealed  point 
is  broken  under  the  water.  As  the 
fountain  is  somewhat  slow  in  start- 
ing, a  small  bulb  of  water,  sealed, 
may  be  placed  within  the  flask, 
and,  when  the  apparatus  is  to  be 
used,  can  be  broken  by  a  quick 
shake.  The  water  will  absorb  hy- 
drochloric acid,  create  a  partial 
vacuum,  and  the  blue  litmus  solu- 
tion will  run  in. 

32.  DECOMPOSITION  OF  HY- 
DROCHLORIC ACID  INTO  ONE  VOL- 
UME OF  CHLORINE  AND  ONE  OF 
HYDROGEN.  —  The  apparatus  to 


Fig.  28. 


*  A  burette  is  a  graduated  glass  tube  about  500  m.m.  in  length,  one  end  of  which  termi- 
nates in  a  narrow  tube,  closed  with  a  glass  stopcock;  a  measured  quantity  of  liquid  can  there- 
fore  be  run  out  of  this  instrument  by  opening  the  stopcock  at  the  bottom. 


LABORATORY  NOTES. 


513 


be  used  is  shown  by  Fig.  29;  pure,  dry  hydrochloric  acid  gas  is  introduced 
through  the  glass  stopcock  of  the  apparatus  (which  has  previously  been  filled 
with  dry  mercury)  until  one  arm  of  the  U-shaped  tube  is  about  two-thirds  full, 
the  mercury  being  allowed  to  run  off  at  the  lower  tap  as  fast  as  hydrochloric  acid 
enters  through  the  upper  stopcock;  when  enough  gas  has 
been  admitted,  the  upper  and  lower  stopcocks  are  closed, 
the  mercury  being  placed  at  the  same  level  in  both  arms. 
The  hydrochloric  acid  gas  is  now  under  atmospheric  pres- 
sure; the  level  of  the  mercury  is  therefore  marked  by  a 
rubber  ring.  Now  sodium  amalgam  *  is  poured  into  the 
open  arm  until  the  arm  is  quite  full,  a  rubber  stopper  is 
inserted  tightly,  and  then  the  apparatus  is  shaken  from 
side  to  side,  so  that  the  sodium  amalgam  comes  in  con- 
tact with  the  gas;  the  hydrochloric  acid  is  decomposed 
after  the  lapse  of  about  one  minute.  The  gas  which 
remains  is  now  carefully  brought  back  into  the  arm  of  the 
apparatus  which  it  originally  occupied,  the  rubber  stop- 
per removed,  and  the  mercury  run  out  through  the  lower 
tap  until  on  the  same  level  in  both  arms;  it  will  then  be 
seen  that  the  volume  of  hydrogen  left  is  exactly  one-half 
of  that  which  was  occupied  by  the  hydrochloric  acid. 
Take  care  to  have  the  apparatus,  sodium  amalgam,  and 
hydrochloric  acid  perfectly  dry  before  beginning  the 
experiment;  for  sodium  amalgam,  acting  on  water,  will 
generate  hydrogen.  Fi9-  2g- 

33.  THE  PREP- 
ARATION OF  BRO- 
MINE. —  Take  a 
tubulated  and 
glass-stoppered  re- 
tort (  Fig.  30)  of  200 
c.c.  capacity;  place 
in  this  10  grams  of 
manganese  dioxide 
and  20  grams  of 
potassium  bromide 
(these  constituents 
previously  mixed 
in  a  mortar),  clamp 
the  retort  to  a  re- 
tort stand, as  shown 

*  To  prepare  sodium  amalgam,  take  500  grams  of  dry  mercury,  place  in  a  large  clay  cru- 
cible, and  cover  with  an  iron  dish;  now  cut  7  grams  of  clean  sodium  into  pieces  the  size  of  a 
hickory  nut,  and  throw  these  into  the  crucible;  to  start  the  reaction,  heat  a  little  mercury  in 
a  test-tube,  raise  the  cover  on  the  crucible  a  little,  and  pour  in  the  hot  mercury;  an  instan- 
taneous reaction,  accompanied  by  a  flash  of  light,  will  occur;  during  this,  stand  aside  so  as 
not  to  inhale  fumes  of  mercury;  after  all  is  over,  stir  the  sodium  amalgam  with  an  iron  wire; 
place,  when  cool,  in  a  wide-mouthed  glass-stoppered  bottle,  and  keep  for  future  use.  Pre- 
pare sodium  amalgam  under  the  hood. 


Fig  30. 


514 


APPENDIX   OF 


in  the  figure,  and  thrust  the  neck  of  the  retort  far  into  the  neck  of  a  receiver  of 
500  c.c.  capacity.  This  receiver  is  cooled  by  means  of  a  stream  of  water,  the 
escape  of  which  is  provided  for  by  a  battery  jar  with  a  syphon  placed  under  the  re- 
ceiver; the  syphon  should  be  connected  with  a  sink.  When  all  is  ready,  add 
sulphuric  acid  (one  part  sulphuric  acid  to  one  part  of  water)  to  the  manganese 
dioxide  and  potassium  bromide,  put  in  the  glass  stopper  of  the  retort,  and  seal 
the  latter  tightly  by  means  of  a  little  plaster  of  paris  and  water;  allow  the 
retort  to  stand  for  ten  minutes,  and  then  heat  gently;  the  bromine  will  distil 
and  collect  in  the  receiver.  The  operation  should  be  performed  under  the 
hood!  Make  a  solution  of  bromine  in  water,  and  try  its  bleaching  power  as 
you  did  with  chlorine;  fill  a  tube  about  500  m.m.  in  length,  closed  at  one  end, 
with  chlorine  water,  and  invert  over  a  beaker  glass  which  is  partly  filled  with 
the  same  liquid,  and,  after  properly  supporting  the  tube,  place  the  whole  in 
the  sunlight;  do  the  same  with  bromine  water,  placing  the  apparatus  filled 
with  the  latter  beside  that  containing  chlorine  water,  and,  after  24  hours, 
notice  the  amount  of  oxygen  separated  by  each  halogen  (see  page  79). 


Fig.  31. 

34.  PREPARATION  OF  HYDROBROMIC  ACID.  —  To  a  few  particles  of  so- 
dium or  potassium  bromide,  in  a  test-tube,  add  a  few  drops  of  concentrated 
sulphuric  acid,  and  warm  very  gently;  note  color  and  odor  of  the  gas  which  is 
passed  off.  Pure  hydrobromic  acid  cannot  be  prepared  in  this  way.  The 
apparatus  for  the  preparation  of  hydrobromic  acid  for  laboratory  use  is  shown 
by  Fig.  31.  The  flask  A,  which  has  a  tube  fused  into  the  side  of  the  neck,* 
*  A  so-called  fractional  distilling-flask. 


LABORATORY  NOTES.  515 

has  a  capacity  of  300  c.c.  It  is  fitted  with  a  single  bored  rubber  stopper,  into 
which  is  inserted  a  drop  funnel.  The  side  tube  connects  with  a  glass  tube,  B, 
which  is  about  300  m.m.  in  length,  and  which  is  filled  with  pieces  of  brick  which 
have  been  moistened  and  rolled  in  red  phosphorus ;  the  upward  slant  to  this 
tube  is  necessary  to  prevent  pieces  of  phosphorus  and  impure  water  from 
being  carried  over  into  the  funnel  tube,  C.  The  latter  is  interposed  between  the 
tube  B  and  the  water  which  is  in  the  beaker  glass,  solely  to  prevent  the  latter 
from  "sucking  back"  as  soon  as  the  current  of  hydrobromic  acid  becomes  so 
feeble  that  solution  of  the  gas  in  water  takes  place  so  rapidly  that  the  gas, 
which  is  being  generated,  is  unable  to  keep  the  liquid  out  of  the  apparatus. 
By  reason  of  the  interposition  of  the  funnel  tube,  the  water  can  be  forced  back 
no  farther  than  the  tube,  provided  its  mouth  is  placed  so  as  to  just  touch  the 
surface  of  the  liquid  in  the  beaker.  Charge  the  generating-flask,  A,  with  25 
grams  of  red  phosphorus,  and  then  add  just  barely  enough  water  to  cover  the 
latter*  fill  the  drop  funnel  above  half  full  of  bromine  (hood !),  connect  all  parts 
of  the  apparatus,  and  allow  the  bromine  to  fall  on  the  phosphorus  slowly,  drop 
by  drop  (care!);  each  drop  of  bromine  will  cause  a  flash  of  light  and  the  for- 
mation of  hydrobromic  acid.  After  all  of  the  water  in  the  flask,  A,  and  the 
tube,  J5,  has  been  saturated  with  hydrobromic  acid,  the  acid  will  escape  into  the 
funnel  tube,  C,  and  can  be  collected  in  a  beaker  of  water,  or,  as  the  dry  gas,  by 
downward  displacement,  as  is  done  with  chlorine.  White  crystals  of  phospho- 
nium  bromide  (see  page  217)  will  ultimately  clog  the  apparatus  if  too  little 
water  is  present.  This  may  become  so  serious  as  entirely  to  prevent  the  flow  of 
gas  through  B,  in  which  case  an  explosion  will  inevitably  result.  Prevent  this 
accident  by  adding  a  few  drops  of  water  to  the  generator  A  as  soon  as  you  see 
phosphonium  bromide  crystals  forming.  Perform  the  same  experiments  with 
the  solution  of  hydrobromic  acid  as  you  did  with  hydrochloric  acid.  Fill  one 
jar  with  dry  hydrobromic  acid  gas,  and  invert  a  jar  of  chlorine  of  the  same 
size  over  it;  observe  the  separation  of  bromine  and  the  rate  of  diffusion  of 
the  gases. 

35.  PREPARATION  OF  IODINE.  —  Perform  this  operation  with  the  same 
apparatus  and  same  proportions  as  you  used  in  the  preparation  of  bromine; 
you  can  omit  the  cooling  of  the  receiver  ;  of  course,  substitute  potassium 
iodide  or  sodium  iodide  for  the  bromide  which  you  used  in  Note  33. 

36.  PREPARATION  OF  HYDROIODIC  ACID.  —  The  apparatus  is  the  same 
as  that  used  for  the  preparation  of  hydrobromic  acid  (Fig.  31),  with  the  ex- 
ception that  the  bricks,  covered  with  moist  red  phosphorus,  can  be  omitted 
from  the  tube  J?,  as  all  of  the  iodine  which  may  pass  over  from  the  generating 
flask  A  will  be  condensed  by  the  cold  glass  tube.     In  charging  the  apparatus, 
place  50  grams  of  iodine  in  the  generating  flask  A,  add  to  this  10  grams  of 
water  ;  replace  the  drop  funnel  used  in  the  hydrobromic  experiment  by  an 
ordinary  funnel,  which  can  be  stoppered  by  a  glass  rod  thrust  into  the  neck 
and  ground  into  the  tip  so  as  to  be  water-tight  ;  through  this  funnel,  after  all 

*  In  preparing  hydrobromic  acid  the  greatest  care  must  be  taken  not  to  add  too  much 
water,  otherwise  the  liquid  will  dissolve  the  hydrobromic  acid  as  fast  as  it  is  generated. 


516 


APPENDIX   OF 


the  connections  of  the  apparatus  are  tight,  gradually  drop  upon  the  iodine 
red  phosphorus,  which  has  been  stirred  with  water  to  a  thick  paste.  By 
following  these  directions  the  generation  of  phosphonium  iodide,  which 
invariably  results  if  the  apparatus  is  arranged  exactly  as  in  the  preparation 
of  hydrobroinic  acid,  is  avoided.  Care  must  be  taken  not  to  add  the  phos- 
phorus too  rapidly,  otherwise  an  explosion  would  result.  The  proportions 
which  are  most  successful  for  the  preparation  of  hydroiodic  acid  are  :  10  parts 
of  iodine,  5  parts  of  phosphorus,  and  2  parts  of  water. 

37.     THE  PREPARATION  OF  SULPHUR.  —  The  formation  of  sulphur  by 
the  action  of  hydrogen  sulphide  on  sulphur  dioxide.     The  apparatus  is  shown 


Fig.  32. 

by  Fig.  32.  The  flask  C,  with  two  lateral  tubulures,  has  500  c.c.  capacity  ;  in 
the  flask,  A  place  about  10  grams,  of  copper  shavings,  stopper  with  a  double 
bored  rubber  stopper  into  which  is  inserted  a  safety  tube  and  delivery  tube, 
add  100  grams  of  concentrated  sulphuric  acid  through  the  safety  tube,  and 
connect  with  C,  as  shown  in  the  cut.  The  double-necked  (Woulff)  wash- 
bottle  -B,  which  contains  a  little  water,  is  connected  with  a  hydrogen  sulphide 
generator.  The  latter  is  the  same  as  that  used  for  hydrogen  (Note  12),  with 
the  exception  that  the  zinc  is  replaced  by  10  grams  of  ferrous  sulphide  ;  the 
flask  C  has  its  mouth  emptying  in  a  beaker,  containing  a  solution  of  sodium 
hydroxide  for  the  purpose  of  absorbing  the  excess  of  gases.  When  all  is 


LABORATORY   NOTES. 


517 


ready,  connect  all  parts  of  the  apparatus,  and  heat  the  copper  and  sulphuric 
acid  in  A,  until  the  ebullition  indicates  that  the  formation  of  sulphur  dioxide 
has  begun,  then  remove  the  flame ;  next  add  dilute  sulphuric  acid  to  the  ferrous 
sulphide  in  the  generating-flask  connected  with  B.  The  sulphur  dioxide 
and  hydrogen  sulphide  will  meet  in  C  ;  after  a  time,  small  drops  of  plastic 
sulphur  will  collect ;  the  latter  soon  become  opaque  and  yellow,  changing  into 
ordinary  rhombic  sulphur.* 

38.  THE  DISTILLATION  OF  SULPHUR.  —  Bend  a  test-tube,  and  clamp  it 
to  a  retort  stand,  as  is  shown  by  Fig  33.     Place  3  or  4  grams  of  sulphur  in  the 
test-tube  and  heat;  collect  the  vapors  in  a 

beaker  of  cold  water.  Endeavor  to  ob- 
serve all  of  the  phenomena  mentioned  on 
page  92,  and  observe  the  formation  of 
flowers  of  sulphur  on  the  surface  of  the 
cold  water. 

39.  THE  CRYSTALLIZATION  OF  SUL- 
PHUR. —  Rhombic    sulphur.       Take    some 
dry  carbon  bisulphide,  t  and  place  it  in  a 
test-tube ;  add  roll  sulphur  to  this  and  then 
stopper;  allow  to  stand  until  the  carbon 
bisulphide  has  taken  up  as  much  sulphur 
as  it  will ;  now  pour  the  clear  solution  into 
a  second  clean  test-tube,  stopper  the  latter 
with  some  cotton,  and  put  aside  in  a  quiet 
place.      After    a    time,    fine,    transparent 
crystals  of  rhombic  sulphur  will  separate. 

Monoclinic  sulphur.  Melt  100  grttms  of  sulphur  in  a  flat  porcelain  evaporat- 
ing-dish,  and  then  allow  to  cool  until  a  crust  has  formed  over  the  surface; 
perforate  this  crust  by  means  of  a  glass  rod  before  the  entire  mass  becomes 
solid,  and  then  pour  off  the  sulphur  which  has  not  solidified,  through  the  hole 
which  has  been  formed.  The  bottom  of  the  evaporating-dish  will  be  covered 
with  colorless,  transparent  needles  of  monoclinic  sulphur  crystals. 

40.  THE   PREPARATION  OF  HYDROGEN    SULPHIDE.  —  The   apparatus 
used  is  the  same  as  that  employed  to  generate  hydrogen  (Fig.  34).     In  the 
generating-flask  place  20  grams  of  ferrous  sulphide,  broken  to  about  the  size 
of  a  bean;  connect  the  delivery  tube  with  one  of  the  bottles  from  the  train 
used  in  drying  hydrogen  (this  bottle  should  contain  a  little  water,  so  as  to 
retain  any  acid  fumes  which  may  pass  over)  ;  after  all  connections  are  made, 
pour  dilute  sulphuric  acid  through  the  thistle  tube.      Collect  the  sulphuretted 
hydrogen  which  passes  off,  as  you  did  with  chlorine  and  hydrobromic  acid,  by 
displacement  of  air,  in  dry  bottles  (such  as  you  used  for  hydrogen  and  oxygen); 
after  three  bottles  have  been  filled  in  this  way,  pass  the  remainder  of  the  gas 


Fig.  33. 


*  It  takes  very  nearly  an  hour  to  form  a  good  deposit  of  sulphur. 

t  Dry  the  carbon  bisulphide  by  shaking  it  with  fused  calcium  chloride. 


518 


APPENDIX   OF 


into  a  bottle  filled  with  water.     Apply  a  lighted  taper  to  one  of  the  bottles 

filled  with  dry  gas  ;  place  a 
strip  of  filter  paper  which 
is  soaked  with  a  solution  of 
lead  acetate  in  the  second  ; 
pass  chlorine  from  the  small 
chlorine  generator,  which 
you  prepared  for  the  experi- 
ment in  Note  29,  into  the 
third.  Take  the  solution  of 
hydrogen  sulphide  which 
you  have  prepared,  and  add 
a  little  of  it  to  a  solution  of 
copper  sulphate  ;  to  a  solu- 
tion of  cadmium  nitrate;  to 
a  hydrochloric  acid  solution 
of  arsenic  trioxide  ;  to  a 
slightly  acid  solution  of  stan- 
nous  chloride*  (see  page  09 
of  the  text).  Experiments 
with  hydrogen  sulphide 
must  be  conducted  under  a 
hood ! 

41.  THE  DECOMPOSI- 
TION OF  HYDROGEN  SUL- 
PHIDE BY  HEAT.  —  This 

decomposition  can  best  be  shown  by  the  apparatus  depicted  in  Fig.  35. 
This  consists  of  a  flask  of  about  1  litre  capacity,  well 
stoppered  with  a  double  bored  rubber  stopper  ;  into 
this  two  glass  tubes  are  fitted,  through  which  run  two 
pieces  of  tolerably  thick  copper  wire,  the  tips  of  the 
glass  tubes  being  fused  around  these  wires;  the  ex- 
tremities of  the  copper  wires  are  connected  by  a 
platinum  wire  as  shown  in  the  cut.  Fill  the  flask 
completely  with  gaseous  hydrogen  sulphide,  insert  the 
stopper,  and  pass  an  electric  current  through  the  plat- 
inum wire  by  attaching  the  free  ends  of  the  copper 
wires  to  the  two  poles  of  a  battery  (the  current  should 
be  just  sufficient  to  cause  the  wire  to  glow)  ;  the 
hydrogen  sulphide  will  decompose  at  the  line  of  con-  p. 

tact  of  the  hot  wire. 

The  experiments  leading  to  the  preparation  of  the  oxygen  acids  and  salts 
of  chlorine,  bromine,  and  iodine  are  made  sufficiently  clear  in  the  text,  and 
need  give  no  difficulty,  as  the  pupil  is  now  able  to  prepare  the  halogens  and 

*  The  pupil  can  select  solutions  of  the  salts  of  other  metals,  and  study  the  solubility  in 
acids  and  in  alkalies  of  the  sulphides  produced  by  the  addition  of  hydrogen  sulphide. 


Fig.  34. 


LABORATORY   NOTES. 


519 


can  use  them  in  the  formation  of  the  various  compounds  mentioned  in 
chapters  xviii.  and  xix.  In  performing  these  experiments,  he  should  confine 
himself  to  the  preparation  and  reactions  of  those  salts  which  are  formed  by  add- 
ing chlorine,  bromine,  and  iodine  to  solutions  of  potassium  hydroxide,  and  to  the 
formation  of  calcium  hypochlorite  by  passing  chlorine  over  slaked  lime,  and  to 
the  decomposition  of  calcium  hypochlorite  by  hydrochloric  and  sulphuric  acids. 

42.  THE  PREPARATION  OF 
SULPHUR  DIOXIDE.  —  The  appa- 
ratus is  shown  in  Fig.  36.  The 
generating-flask  should  be  of  500 
c.c.  capacity;  in  this  place  about 
twenty  grams  of  copper  shavings, 
connect  all  parts  of  the  apparatus, 
add  about  100  c.c.  of  concentrated 
sulphuric  acid  through  the  safety 
tube,  and  heat  by  means  of  a  Bun- 
sen  burner  ;  when  the  gas  begins 
to  pass  off,  lower  the  flame  so 
as  to  secure  a  regular  evolution. 
The  wash-bottle  contains  concen- 
trated sulphuric  acid.  Collect  three 
jars  of  the  gas  by  displacement  of 
air,  and  pass  the  remaining  gas  into 
water.  See  if  the  gas  will  burn  or 
will  support  combustion;  put  a  moist 
strip  of  colored  calico  into  one  of  the 
jars,  and  a  small  red  rose,*  which 
has  been  moistened,  into  another. 
To  portions  of  the  solution  of  sul- 
phur dioxide  in  water  add,  succes- 
sively, a  dilute  alcoholic  solution  of  Fig.  36. 
iodine ;  a  solution  of  bromine ;  a  solution  of  ferric  chloride  ;  a  solution  of  po- 
tassium dichromate.  (All  of  these  reagents 
will  illustrate  the  reducing  power  of  sulphur 
dioxide.) 

43.  LIQUID  SULPHUR  DIOXIDE.  —  The 
most  convenient  form  in  which  to  use  sulphur 
dioxide  in  the  laboratory  is  as  a  liquid.  The 
apparatus,  shown  by  Fig.  37,  is  placed,  with  its 
stopcocks  open,  in  a  dish  filled  with  a  mixture 
of  pounded  ice  and  salt,t  and  then  a  slow  cur- 
rent of  sulphur  dioxide,  generated  as  explained 
in  Note  42,  is  passed  through.  In  order  to  cool 
the  gas  perfectly  it  should,  after  leaving  the 
drying-flask,  traverse  a  spiral  glass  tube  which 


37. 


*  Roses  are  best  to  use  for  bleaching  with  sulphur  dioxide;  many  other  red  flowers 
bleach  very  slowly,  some  not  at  all. 

t  In  making  a  freezing  mixture,  do  not  spare  the  salt. 


520 


APPENDIX   OF 


is  placed  in  a  jar  and  cooled  with  snow  and  salt.  When  a  sufficient  quantity 
of  the  gas  has  been  liquefied,  all  of  the  stopcocks  are  closed,  and  the  apparatus 
is  set  aside  until  wanted.  When  gaseous  sulphur  dioxide  is  required,  the 
stopcock  c  is  opened,  and  a  portion  of  the  gas  run  into  the  small  bulb;  c  is 
then  closed,  and  by  opening  b  the  gas  can  be  used  without  interfering  with 
the  liquid  in  the  large  bulb. 

44.     THE  MANUFACTURE  OF  SULPHURIC  ACID.  — A  laboratory  apparatus 
for  illustrating  the  manufacture  of  sulphuric  acid  is  shown  in  Fig.  38.     A 

glass  globe  (Y!)  of  about  5 
litres  capacity  is  fitted  with  a 
rubber  stopper  through  which 
five  holes  are  bored.  Three 
small  flasks  (a,  b,  c)  are  con- 
nected with  this  by  means  of 
glass  tubes  extending  to  the 
middle  of  the  large  one,  and 
through  the  remaining  holes 
two  tubes,  also  extending  to 
the  middle  of  the  large  flask, 
are  fitted.  Sulphur  dioxide 
is  generated  in  a  by  heating 
copper  shavings  and  sulphu- 
ric acid,  nitric  oxide  in  b  by 
Fig'  38'  means  of  copper  and  dilute 

nitric  acid  (no  heat  is  required),  and  steam  is  supplied,  as  wanted,  by  heating 
water  in  c.  Air  can  be  forced  in  at  d  by  connecting  the  tube  with  a  bellows, 
and  the  tube  e,  connected  with  a  hood,  is  left  as  a  vent-hole  for  the  escape  of 
gases.  When  nitric  oxide  comes  in  contact  with  air,  it  is  oxidized  to  a  mix- 
ture of  NO2  and  N2  O3  ;  so  that,  after  that  change  has  taken  place,  we  have 
the  gas  present  which  is  necessary  for  the  formation  of  sulphuric  acid  from 
sulphurous  acid.  If  very  little  steam  is  admitted,  we  can  easily  see  the  forma- 
tion of  nitrosyl-sulphuric  acid;  for  the  large  globe  becomes  covered  with  frost- 
like  crystals  of  that  substance.  If  an  excess  of  steam  is  admitted,  these 
crystals  disappear  and  sulphuric  acid  is  formed,  the  latter  collecting  as  an  oily 
liquid.  By  varying  the  amounts  of  nitric  oxide,  sulphur  dioxide,  air,  and 
steam,  we  can  study  all  of  the  phases  of  sulphuric  acid  manufacture. 

EXPERIMENTS  WITH  SULPHURIC  ACID.  —  Take  some  concentrated  sulphu-/ 
ric  acid,  and  add  it  to  a  little  cane  sugar  which  you  have  placed  in  a  test-tube; 
stir  the  mixture  with  a  glass  rod,  and  allow  to  stand  under  the  hood ;  add  9.8 
grams  of  concentrated  sulphuric  acid  to  1.8  grams  of  water,  place  the  liquid  in 
a  small  flask,  and  surround  the  latter  with  crushed  ice  and  salt;  crystals  of 
H4 SO3  will  form;  warm  the  H4  SO5  until  melted,  and  once  more  add  1.8  grams 
of  water,  place  in  the  freezing  mixture,  and  crystals  of  H6  SO6  will  separate ; 
melt  the  HG  SO6  which  you  have  made,  and  then  gradually  add  more  water, 
and  notice  if  there  is  any  increase  of  temperature.*  Add  some  dilute  sulphu- 

*  In  this  experiment  it  will  be  necessary  to  add  rooter  to  sulphuric  acid  ;  be  sure  to  hold 
the  flask  at  a  safe  distance  while  pouring  in  the  water,  otherwise  the  heat  which  is  generated 
might  cause  the  water  to  boil  and  spatter  drops  of  sulphuric  acid. 


LABORATORY   NOTES. 


521 


ric  acid  to  a  solution  of  barium  chloride;  to  a  solution  of  strontium  chloride; 

and  to  two  solutions  of  calcium  chloride,  the  first  of  which  is  very  dilute,  the 

second  tolerably  concentrated,  and  note  the  result  (see  page  416). 

45.     To  ISOLATE  NITROGEN  FROM  THE  ATMOSPHERE.  —  The  apparatus  is 

shown  by  Fig.  30.     Take  a  bell  jar  of  3  litres  capacity,  invert  it  over  a  glass 

basin,  or  over  your  pneumatic  trough;  prepare  • 
a  float  made  of  a  flat  cork,  on  one  side  of  which 
you  have  fastened  a  porcelain  crucible  cover  by 
forcing  the  handle  firmly  into  the  cork;  place 
this  cork,  with  the  crucible  cover  up,  in  your 
pneumatic  trough ;  place  a  piece  of  phosphorus 
the  size  of  a  bean  on  the  cover,  and  light  the 
phosphorus  with  a  hot  wire  (care  in  handling 
phosphorus!);  now  invert  the  bell  jar  over  the^ 
float,  and  slightly  raise  the  stopper  at  the  top, 
so  that  the  gas  within* 
which  necessarily  ex- 


Fig.  39. 


pands  very  greatly  (owing  to  the  heat  given  off  by  the 
burning  phosphorus),  can  quietly  escape.  If  this  pre- 
caution is  omitted,  the  air  will  be  forced  out  at  the  bot- 
tom of  the  jar  in  large  bubbles;  the  disturbance  may 
even  tip  over  your  phosphorus  float.  After  the  violent 
combustion  is  over,  insert  the  stopper  of  the  bell  jar  and 
allow  the  gas  to  cool ;  water  will  rise  in  the  jar  to  take 
the  place  of  the  oxygen  which  has  gone  to  form  phos- 
phorus pentoxide,  and  also  to  take  the  place  of  the  air 
which  has  been  expelled.  To  test  the  gas  remaining, 
add  enough  water  to  the  pneumatic  trough  to  make  the 
level  within  and  without  the  bell  jar  alike,  and  then  in- 
troduce a  lighted  taper.* 

46.  THE  COMPOSITION  OF  THE  ATMOSPHERE.  —  A 
crude  method  of  determining  the  volumetric  composi- 
tion of  the  atmosphere  is  by  means  of  the  apparatus 
shown  by  Fig.  40.  Divide  a  long  glass  tube,  closed  at 
one  end,  into  five  equal  parts  by  means  of  rubber  rings. 
Invert  this  over  a  long  cylinder  containing  water  so  that 
the  level  without  and  within  is  at  the  first  ring,  and  then 
clamp  the  tube  in  place,  Now  fix  a  piece  of  phosphorus 
on  a  long,  sharp-pointed  copper  wire,  taking  care  not  to 
touch  the  phosphorus  with  the  hands,  bend  the  wire  as 
shown  in  the  cut,  thrust  the  phosphorus  up  into  the 
tube,  and  set  the  apparatus  aside  for  two  days.  The 
oxygen  of  the  enclosed  air  will  then  be  entirely  ab-  Flg"  40" 

sorbed;  and,  by  sinking  the  tube  so  that  the  level  of  the  water  without  and 
within  is  the  same,  the  amount  of  nitrogen  in  the  air  can  be  ascertained.     By 

*  The  sides  of  the  bell  jar  become  coated  with  a  red  amorphous  solid  during  this  opera- 
tion; this  substance  is  a  sub-oxide  of  phosphorus  with  the  probable  formula  P4O;  it  is  not 
red  phosphorus,  as  is  generally  supposed. 


522  APPENDIX   OF 

noting  the  height  of  the  barometer  before  and  after  the  experiment,  and  then 
applying  the  necessary  corrections,  quite  accurate  results  can  be  obtained ;  but, 
if  such  are  required,  a  carefully  graduated  tube  must  be  substituted  for  the 
crudely  divided  one  indicated. 

In  order  to  measure  accurately  the  relative  amounts  of  oxygen  and  nitro- 
gen in  the  atmosphere,  the  eudiometer  (Fig.  23,  Note  20)  is  employed.  The 
instrument  should  have  a  capacity  of  100  c.c. ;  it  is  partially  filled  with  mer- 
cury, and  inverted  over  the  mercury  trough  so  that  about  25  c.c.  of  air  will 
remain  enclosed ;  about  14  c.c.  of  hydrogen  are  now  run  in,  by  slanting  the 
tube  and  placing  the  delivery  tube  of  a  hydrogen  apparatus  (which  is  gen- 
erating pure  and  dry  hydrogen)  under  its  mouth;  take  all  of  the  precautions 
mentioned  in  Note  20,  and  ignite  the  mixture  of  gases  with  an  electric  spark; 
be  sure  to  read  accurately  the  volume  of  air  and  the  volume  of  hydrogen 
before  the  explosion,  and  also  to  measure  the  height  of  the  column  of  mercury 
as  indicated  in  Note  20;  after  the  explosion  read  the  volume  of  remaining  gas 
and  reduce  to  standard  Conditions,  exactly  as*was  done  before.  The  hydrogen 
will  have  united  with  the  oxygen  to  form  water;  therefore,  one-third  of  the 
volume,  by  which  the  mixture  of  the  gases  which  were  enclosed  in  the  eudio- 
meter has  diminished,  must  have  been  oxygen.* 

47.  THE  PRESENCE  OF  CARBON  DIOXIDE  IN  THE  ATMOSPHERE.  — 
Take  one  of  the  train  of  wash-bottles  which  was  used  in  drying  hydrogen, 
clean  it,  and  fill  it  with  clear  lime-water,  and  attach  to  a  Bunsen  aspirator 
which  you  have  fastened  to  the  hydrant,!  and  the  suction  tube  of  which  is 
connected  with  your  wash-bottle  in  such  a  way  that  air  will  be  drawn  through 
the  apparatus  in  the  same  direction  as  hydrogen  is  forced  through  it,  as  indi- 
cated in  Fig.  18.     The  lime-water  will   soon  become   turbid,  owing  to  the 
formation  of  calcium  carbonate ;  if  you  use  two  wash-bottles,  each  of  which 
contains  clear  lime-water,  then  the  one  into  which  the  air  first  passes  will 
become  turbid,  while  the  second  will  remain  clear.     The  presence  of  moisture 
in  the  atmosphere  can  be  shown  by  exposing  an  open  beaker,  containing  con- 
centrated sulphuric  acid,  to   the   air;   after  some   days  the  volume  will  be 
observed  to  have  increased,  while  the  acid  has  become  diluted  with  water. 

48.  THE  PREPARATION  OF  AMMONIA.  —  The  apparatus  for  the  prepara- 

*  If  the  temperature  and  barometer  are  the  same  before  and  after  the  experiment,  no 
correction  need  be  made  excepting  for  the  difference  caused  by  the  change  in  the  pressure 
on  the  enclosed  gas,  brought  about  by  the  differing  height  of  the  column  of  mercury  in  the 
eudiometer;  the  latter  must  be  carefully  noted,  1st,  when  air  alone  is  in  the  eudiometer,  2d, 
when  hydrogen  has  been  admitted,  3d,  after  the  explosion.  As  the  air  contains  an  unknown 
amount  of  moisture,  saturate  it  with  water  vapor  by  admitting  a  drop  of  water  above  the 
mercury  in  the  eudiometer;  and  then,  unless  temperature  and  barometric  pressure  are  differ- 
ent after  the  experiment  than  they  were  before,  there  will  be  no  necessity  for  paying  atten- 
tion to  the  amount  of  water  present,  for  the  gases  will  be  saturated  with  water  before  and 
after  the  experiment  (see  page  173). 

t  A  Bunsen  glass  aspirator  is  a  cheap  instrument  which  aspirates  air  by  using  the  water 
pressure  of  a  hydrant;  it  is  indispensable  for  laboratory  work,  and  should  be  kept  in  the 
desk,  ready  for  use.  It  is  attached  by  means  of  rubber  tubing,  and  its  mechanism  will 
explain  itself  when  the  instrument  is  handled.  A  more  effective,  but  also  more  expensive, 
instrument  is  a  Chapman  brass  aspirator;  this  can  also  be  fastened  to  the  hydrant;  both 
forms  of  aspirator  can  be  purchased  of  any  instrument  dealer. 


LABORATORY  NOTES. 


523 


tion  of  ammonia  is  shown  by  Fig.  41.     The  generating-flask  is   of  500  c.c. 

capacity;  in  it  are 
placed  50  grams  of 
ammonium  chloride 
and  100  grams  of 
slaked  lime  (pre- 
pare slaked  lime  by 
slowly  pouring 
water  on  quick-lime 
until  the  latter  fi- 
nally crumbles  to  a 
powder) ;  connect 
the  generating-flask 
with  a  drying-cylin- 
der, the  latter  being 
filled  with  small 
pieces  of  quick- 
1  i  m  e,  which  are 
present  for  the  pur- 
pose of  drying  the 
gas;  collect  ammo- 
Fig.  41.  nia  over  mercury  or 

by  displacement  of 

air  in  a  jar  which  is  held  mouth  downward,  for  ammonia  is  specifically  lighter 
than  air.  When  all  connections  are  made,  add  enough  water  to  the  mixture 
of  ammonium  chloride  and  slaked  lime  in  the  flask  to  cause  the  latter  to  roll 
into  lumps  on  shaking;  now  heat  gently  and  ammonia  will  pass  off.  The 
drying-cylinder  is  a  so-called  "  Fresenius  "  drying-tower,  at  the  bottom  open- 
ing of  which  the  gas  enters;  at  the  top,  after  traversing  the  intermediate  space 
filled  with  pieces  of  quick-lime,  it  escapes.  These  drying-towers  are  very 
convenient  for  laboratory  use,  and,  if  possible,  should  be  kept  on  hand. 
Ammonia  cannot  be  dried  over  calcium  chloride,  because  it  combines  with  that 
substance ;  it  is  self-evident  that  the  gas  cannot  be  dried  by  sulphuric  acid.  • 

49.  EXPERIMENTS  WITH  AMMONIA.  —  The  combustion  of  ammonia  may  be 
shown  by  filling  a  large  test-tube  with  oxygen,  then  passing  in  some  ammonia 
gas,  so  as  partially  to  displace  the  oxygen,  and  then  quickly  approaching  the 
mouth  of  the  test-tube  to  a  gas  flame ;  a  weak  explosion  will  follow.*  Pass 
some  of  the  ammonia,  which  was  generated  by  the  experiment  mentioned  in 
Note  48,  into  water,  and,  with  the  solution  of  ammonia  so  produced,  perform 
the  following  experiments  :  Place  some  ammonia  in  a  beaker  glass;  add  a 
few  drops  of  blue  litmus  solution,  and  then  carefully  add  hydrochloric  acid 
from  a  burette,  until  the  solution  is  neutral  (see  Note  30).  Pour  the  liquid  in 
the  beaker  into  an  evaporating-dish,  and  evaporate  to  dryness  on  a  water  bath 
(Note  21);  heat  a  little  of  the  salt  which  separates,  on  a  piece  of  platinum 


*  Take  care  to  wrap  a  towel  around  the  test-tube  before  bringing  its  mouth  to  the 
flame. 


524 


APPENDIX    OF 


Fig.  42. 


foil;  heat  some  of  the  salt  with  slaked  lime  and  water  in  a  test-tube;  repeat 
the  same  experiments,  substituting  nitric  acid  and  sulphuric  acid  for  hydro- 
chloric acid. 

The  absorption  of  ammonia  by  charcoal  can  be  demonstrated  by  the  appara- 
tus shown  by  Fig.  42.  This  is  a  test-tube  filled  with  dry  ammonia  and 
inverted  over  a  basin  of  mercury;  a  small  piece  of 
charcoal,  which  has  previously  been  glowed  out  in  a 
Bunsen  burner,  is  introduced,  and  the  test-tube  is  then 
clamped  with  its  mouth  under  the  surface  of  the  mer- 
cury; the  ammonia  will  be  absorbed,  and  the  mercury 
will  rise  in  the  tube. 

50.  THE  SOLUBILITY  OF  AMMONIA  IN  WATER.  — 
Use'  the    same    apparatus   which    you    employed    for 
demonstrating  the  solubility  of  hydrochloric  acid  (Note 
31,  Fig.  28),  filling  the  flask  with  dry  ammonia  gas  by 
displacement  of  air. 

51.  AMMONIUM  AM  A  La  AM.  —  In  a  narrow  cylinder 
of  225  c.c.  capacity,  place  20  grams  of  ammonium  chlo- 
ride, add  enough  water  so  as  just  to  cover 

the  salt,  and  then  pour  on  sodium  amalgam, 
prepared    as    indicated   in   the  foot-note  to 

Note  32:  the  ammonium  amalgam  will  begin  to  form  at  once,  and 
the  operation  can  be  hastened  by  gently  stirring  with  a  glass  rod. 

52.      THE    PREPARATION    OF     NITROUS 
OXIDE.  —  The   apparatus  which  it  is  best  to 
use  for  the   preparation  of  this  gas  is  shown 
by  Fig.  43.     This  consists  of  a  300  c.c.  flask, 
stoppered  with  a  single  bored  rubber  stopper 
which  is   connected  with  a   delivery  tube  and 
safety  bottle.    In  this  flask,  place  10  to  15  grams 
of    crystallized    ammonium 
nitrate,*  and    heat  to  a  tem- 
perature   just   sufficiently   high 
to  cause  a  regular  flow  of  the 
gas ;  collect  over  the  pneumatic 
trough     by     displacement     of 
water.      Introduce     a     lighted 
taper   into    a    jar    filled   with 
nitrous  oxide,  and  see  whether 
it  burns  or  supports   combus-  Fig.  43. 

tion;    repeat   the    experiments 

given  in  Notes  6  and  7,  using  nitrous  oxide  instead  of  oxygen ;  inhale  a  little 
of  the  gas. 

*  Test  a  little  of  your  ammonium  nitrate  by  heating  in  a  test-tube  before  you  proceed  to 
the  decomposition  of  larger  quantities  of  the  salt. 


0 


0 


LABORATORY   NOTES. 


525 


53.  PREPARATION  OF  NITRIC  OXIDE. —  Use  the  apparatus  shown  by 
Fig.  44,  charge  the  flask  with  20  grams  of  copper  shavings,  cover  with  water, 
and  slowly  add  ordinary  nitric  acid,  waiting  for  the  reaction  to  begin  after 

each  addition;  after  all  the 
brown  fumes  which  are  at 
first  developed  have  disap- 
peared, collect  the  gas  over 
the  pneumatic  trough  by  dis- 
placement of  water.  Do  not 
heat. 

54.  EXPERIMENTS  WITH 
NITRIC  OXIDE.  —  Take  one 
of  the  bottles  which  you  have 
filled  with  nitric  oxide  and 
turn  it  mouth  upward;  the 
brown  fumes  of  the  higher 
oxides  of  nitrogen  will  ap- 
pear where  contact  with  the 
air  takes  place;  repeat  the 
experiments  illustrating  com- 
bustion which  you  performed 
with  nitrous  oxide  and  with 
oxygen,  and  satisfy  yourself 
as  to  whether  nitric  oxide 
supports  combustion  as  read- 
ily afis  either  of  those  two 
gases;  place  a  few  drops  of 
carbon  bisulphide  in  one  of  your  bottles  of  nitric  oxide,  after  you  have  re- 
moved it  from  the  pneumatic  trough,  and  covered  its  mouth  with  a  piece  of 
glass  plate,  taking  care  not  to 
admit  any  more  air  than  is  ab- 
solutely necessary  while  pouring 
in  the  carbon  bisulphide;  now 
shake  the  bottle  back  and  forth 
three  or  four  times,  keeping  it 
closed  ;  carefully  allow  the  ex- 
cess of  carbon  bisulphide  to  leak 
out  at  the  place  where  you  hold 
the  glass  cover  on  the  bottle,  and 
then,  while  quickly  removing  the 
cover,  approach  the  mouth  of 
the  bottle  to  a  gas  flame.  (Pages 
200  and  201.) 

55.  PREPARATION  OF  NI- 
TROGEN PEROXIDE  FROM  NI- 
TRIC OXIDE.  —  The  apparatus  Fig'  45' 


Fig.  44. 


526  APPENDIX   OF 

is  shown  by  Fig.  45.  Fill  with  nitric  oxide  a  glass  flask  of  500  c.c.  capacity, 
tubulated  at  one  side  and  having  a  long  neck ;  invert  the  flask  over  a  basin 
of  water  so  that  its  mouth  is  under  the  liquid ;  connect  the  tube  which  is  fitted 
to  the  side  tubulure  with  an  oxygen  gasometer  (Note  14),  and,  by  opening  the 
wire  pinchcock  which  closes  the  rubber  tube,  admit  a  little  oxygen.  Brown 
fumes  of  nitrogen  peroxide  will  at  once  appear;  the  latter  gas  is,  however, 
rapidly  absorbed  by  the  water  in  the  neck  of  the  flask,  the  following  reaction 
taking  place :  — 

3  NO2  +  H2  O  =  2  HNO3  +  NO  (see  page  198). 

After  the  gas  in  the  glass  globe  has  become  colorless,  owing  to  the  disap- 
pearance of  the  peroxide  and  the  regeneration  of  a  portion  of  the  nitric 
oxide,  add  a  little  more  oxygen,  wait  for  the  absorption  of  the  brown  fumes 
again,  and  repeat  the  experiment  until  all  of  the  nitric  oxide  has  been  used 
up;  the  water  from  the  trough  will  then  have  completely  filled  the  globe.* 

The  preparation  of  nitrogen  peroxide  by  heating  lead  nitrate.  —  The  appa- 
ratus is  the  same  as  that  used  for  the  preparation  of  oxygen  by  heating  mercuric 
oxide  (Note  1,  Fig.  1),  excepting  that  the  gas  must  be  collected  by  displacement 
of  air,  and  not  over  water.  Fill  the  hard  glass  tube  one-quarter  full  of  a  mixture 
of  equal  parts  of  sand  and  lead  nitrate  (the  latter  ground  fine  in  a  mortar), 
pound  the  side  of  the  tube  sharply  on  the  desk  so  as  to  form  a  canal  for  the 
escape  of  the  gas  (Note  3),  and  heat  in  the  combustion  furnace,  taking  care 
to  decompose  the  lead  nitrate  at  the  rear  end  of  the  tube  first;  and,  when  this 
has  yielded  all  of  the  gas  which  it  is  capable  of  doing,  advance  the  flame  grad- 
ually toward  the  mouth  of  the  tube.  By  passing  the  nitrogen  peroxide  through 
the  apparatus  used  in  condensing  sulphur  dioxide  (Note  43,  Fig.  37),  it  can  be 
obtained  as  a  straw-colored  liquid.  Nitrogen  trioxide  can  be  liquefied  by  the 
same  means,  the  fluid  being  indigo  blue  in  color.  Prepare  the  trioxide  by  heating 
a  mixture  of  ordinary  nitric  acid  and  a  little  arsenious  oxide  t  in  the  apparatus 
which  you  used  for  the  preparation  of  sulphur  dioxide,  after  removing  the 
wash-bottle  and  replacing  the  latter  by  the  gas-condensing  apparatus,  which  is 
well  cooled  by  means  of  pounded  ice  and  salt. 

56.  THE  PREPARATION  OF  NITRIC  ANHYDRIDE.  —  The  preparation  of 
this  substance  is  not  infrequently  attended  with  danger;  the  reaction  leading  to 
its  formation  is,  therefore,  scarcely  to  be  attempted  either  in  the  laboratory 
or  in  the  lecture-room.  Place  concentrated  nitric  acid  in  a  tubulated  retort  of 
about  500  c.c.  capacity,  and  then  gradually  add  phosphoric  anhydride  until  the 
mixture  of  that  solid  with  the  nitric  acid  has  formed  a  jelly-like  mass;  cool 
the  retort  during  this  operation  so  that  the  temperature  never  exceeds  0°;  now 
heat  very  gently  on  the  water-bath,  never  allowing  the  temperature  to  exceed 
60°;  distil  the  liquid  which  passes  off  into  a  receiver  (arranged  as  in  the 

*  Of  course,  if  a  globe  which  is  tubulated  is  not  at  hand,  an  ordinary  flask  can  be  used; 
the  oxygen  can  then  be  run  in  under  the  water  by  using  a  bent  delivery  tube. 

t  If  you  have  none  but  powdered  arsenic  trioxide,  be  careful  to  add  it  very  gradually  to 
the  nitric  acid,  and  heat  very  gently,  as  the  reduction  of  the  nitric  acid  is  apt  to  become 
quite  violent;  if  the  porcelain-like  variety  of  arsenious  oxide  is  at  hand,  it  is  much  better  to 
use  that  form. 


LABORATORY   NOTES. 


• 

\ 


••' 


527 

preparation  of  bromine,  Fig.  30),  which  is  kept  cool  by  means  of  ice  and  salt; 
the  nitric  anhydride  will  then  solidify.  Do  not  keep  nitric  anhydride  for  any 
length  of  time  ! 

57.  THE  PREPARATION  OF  NITRIC  ACID.  —  By  passing  electric  sparks 
through  moist  air.     Take  a  eudiometer  tube  (Fig.  23,  Note  20),  stopper  it  with 
a  rubber  stopper,  connect  its  platinum  wires  with  a  battery  and  induction  coil, 
and  allow  electric  sparks  to  pass  through  for  about  an  hour;  the  tube  will 
then  be  seen  to  be  filled  with  brown  fumes  if  it  is  held  against  a  white  back- 
ground, and  a  little  blue  litmus  solution  introduced  into  the  tube  will  be 
turned  red.* 

58.  THE  SAME.  —  By  heating  sodium  nitrate  and  sulphuric  acid.      The 
apparatus  to  be  used  is  shown  by  Fig.  46.     It  is  identical  with  that  employed 


Fig.  46. 


for  the  preparation  of  bromine  (Fig.  30). t  In  the  retort,  place  50  grams  of 
sodium  nitrate,  make  all  connections,  and  then  add  fifty  grams  of  concen- 
trated sulphuric  acid;  warm  gently  until  drops  of  liquid  begin  to  pass  over, 
and  then  endeavor  to  keep  the  retort  about  the  temperature  of  distillation. 
After  a  time  a  crystalline  remainder  will  form  in  the  retort;  wash  this  out, 
evaporate  in  an  open  dish,  and  investigate  the  nature  of  the  crystals  by  heat- 
ing in  a  test-tube;  if  they  consist  of  the  primary  sulphate,  they  will  separate 
sulphuric  acid  on  heating;  repeat  the  experiment,  using  25  grams  of  sulphuric 
acid  to  50  grams  of  potassium  nitrate ;  heat  until  no  more  nitric  acid  passes 
off;  and  then,  after  cooling  and  re-crystallizing  the  remainder,  test  as  you  did 
the  primary  sulphate  (the  secondary  sulphate  will  not  liberate  sulphuric  acid 
on  heating  in  a  test-tube). 

*  Be  careful  to  have  the  blue  litmus  solution  as  nearly  neutral  as  possible,  and  be  sure 
to  add  no  more  than  one  or  two  drops. 

t  Nitric  acid  attacks  rubber;  it  is  therefore  essential  to  have  some  high  melting  paraf. 
fin  in  the  laboratory;  melt  a  little  of  this,  and  coat  the  rubber  stopper  by  dipping  it  into  the 
liquid;  unless  the  temperature  of  any  reaction  becomes  high  enough  to  melt  the  paraffin, 
the  latter  will  afford  a  perfect  protection  to  the  rubber.  A  good  method  is  to  use  a  glass- 
stoppered  retort,  and  to  lute  the  stopper  with  plaster  of  paris. 


528 


APPENDIX    OF 


59.  EXPERIMENTS  WITH  NITRIC  ACID.  —  Make  a  solution  of  indigo  by 
dissolving  a  little  indigo  in  concentrated  sulphuric  acid,  warming  slightly, 
and  then  diluting  with  water  ;  to  this  solution  add  nitric  acid  until  it  is 
bleached  ;  take  a  piece  of  white  silk  ribbon,  and  dip  it  into  tolerably  con- 
centrated nitric  acid;  wash  with  clean  water  and  put  aside;  after  a  time 
examine  its  color  and  texture.  Prepare  fuming  nitric  acid  by  placing 
100  grams  of  ordinary  nitric  acid  in  the  apparatus  used  for  preparing 
bromine  (Fig.  46),  slowly  adding  50  grams  of  concentrated  sulphuric  acid 
and  gently  distilling  ;  by  this  means  the  nitric  acid  is  deprived  of  nearly 
all  water  with  which  it  was  mixed  ;  now  clean  the  retort,  put  the  dis- 
tillate back  into  it,  and  then  add  a  few  pieces  of  starch,  connect  the  ap- 
paratus and  slowly  distil  again  ;  the  starch  will  generate  lower  oxides  of 
nitrogen  (N2O3  and  NOo)  while  it  is  itself  being  oxidized,  and  these  lower 
oxides  will  be  dissolved  by  the  nitric  acid  in  the  receiver  ;  fuming  nitric  acid 
is,  therefore,  nitric  acid  which  contains  lower  oxides  of  nitrogen  (see 
page  206).  Place  some  of  the  fuming  nitric  acid  in  a  test-tube,  as  is  shown 
by  Fig.  47,  place  the  test-tube  inside  of  a  beaker,  in  order  to  render  an 
accidental  cracking  harmless,  warm  the  nitric  acid  slightly,  and  then  drop  a 

red-hot  piece  of  charcoal,  which  is 
cut  the  size  of  a  pea,  into  the  acid 
(perform  this  experiment  under  the 
hood!).  Try  the  solubility  of  various 
metals  (iron,  zinc,  copper,  platinum) 
in  nitric  acid,  and  note  the  gases 
which  pass  off.  Take  some  pieces  of 
zinc,  and  dissolve  them  in  very  dilute 
cold  nitric  acid  ;  evaporate  the  re- 
mainder to  dryness,  and  then  see  if 
you  can  discover  the  presence  of  an 
ammonium  salt  (Note  49)  ;  do  the 
same  with  a  piece  of  magnesium 
(page  206).  Prepare  aqua  regia  by 
mixing  one  part  of  nitric  acid  with 
three  parts  of  hydrochloric  acid;  allow 

to  stand,  and  notice  if  the  odor  of  chlorine  is  perceptible  ;  dissolve  a  small 
piece  of  platinum  in  aqua  regia.  To  a  solution  of  ferrous  sulphate,  add  a  solu- 
tion of  potassium  hydroxide,  and  note  the  appearance  of  the  precipitate ;  then 
heat  a  solution  of  ferrous  sulphate  with  nitric  acid,  add  potassium  hydroxide, 
and  note  the  appearance  of  the  precipitate.  (Ferrous  hydroxide  is  precipi- 
tated in  the  first  case,  ferric  hydroxide  in  the  second  ;  nitric  acid  oxidizes  -ous 
compounds  to  -ic  compounds  [prove  this  also  by  adding  the  acid  to  a  solution 
of  sulphur  dioxide  in  water].) 

THE  DECOMPOSITION  OF  THE  NITRATES.  —  The  decomposition  of  the 
nitrate  of  a  heavy  metal  was  illustrated  in  Note  55.  The  nitrates  of  the 
alkalies  decompose  into  the  nitrite  and  oxygen  when  heated.  Take  some 
potassium  nitrate,  place  in  a  hard  glass  test-tube,  and  heat  for  some  time  to  a 


LABORATORY   NOTES.  529 

bright  red  heat;  bubbles  of  oxygen  will  pass  off.  Allow  the  test-tube  to  cool, 
dissolve  the  remainder  in  water,  add  a  little  iodide  of  potassium  solution 
mixed  with  starch  paste  (Note  22),  and  then  a  drop  of  sulphuric  acid:  iodine 
will  at  once  be  liberated.  Do  the  same  with  some  pure  potassium  nitrate 
dissolved  in  water,  and  note  the  difference.  (Nitrous  acid  at  once  liberates 
iodine  from  iodide  of  potassium  because  it  is  a  very  quick  oxidizer,  just  as 
ozone  and  hydrogen  peroxide  are  [page  50] ;  nitric  acid  liberates  iodine  only 
after  a  considerable  interval  of  time. ) 

60.  THE  PREPARATION  OF  RED  PHOSPHORUS.  —  Take  a  piece  of  hard 
glass  tubing,  seal  one  end,  place  a  piece  of  phosphorus  the  size  of  a  pea  in 
the  tube,  and  seal  the  other  end  by  heating  in  a  glass-blower's  flame  and 
allowing  the  sides  to  fall  together;*  place  the  tube  upright  in  an  iron  crucible 
of  its  own  length,  and  then  fill  the  crucible  with  sand.     Heat  by  means  of  a 
triple  burner.     By  this  method  the  glass  tube  will  be  hot  below  and  tolerably 
cool  above.     At  some  point  in  its  length  the  proper  temperature  of  300°  will 
be  reached;  red  phosphorus  will  deposit  at  that  place.     (Be  sure  to  perform 
this  experiment  under  a  hood,  so  that  if  the  tube  should  explode  the  flying 
glass  can  do  no  damage.)     Never  handle  the  glass  tube  unless  it  is  cold. 
Open  the  tube  by  wrapping  a  towel  around  it,  and  exposing  the  tip  of  the  long 
sealed  end  to  the  flame;  after  a  little  air  has  been  admitted  in  this  way,  you 
can  break  open  the  tube.     Never  attempt  to  break  open  a  sealed  glass  tube 
unless  you  have  taken  this  precaution. 

The  low  kindling  temperature  of  ordinary  phosphorus  can  be  shown  by 
dissolving  a  little  in  carbon  bisulphide,  and  then  pouring  a  few  drops  of  this 
solution  on  a  piece  of  filter  paper.  After  the  carbon  bisulphide  has  evapo- 
rated, the  finely  divided  phosphorus  which  remains  on  the  filter  paper  will 
take  fire  spontaneously.  (Throw  the  solution  of  phosphorus  in  carbon 
bisulphide  down  the  sink  as  soon  as  you  are  through  with  it,  and  then  wash 
the  test-tube.  Be  careful  while  handling  phosphorus. ) 

61.  THE  PREPARATION  OF  PHOSPHINE. — The  apparatus  is  shown  by 
Fig.  48.     (See  p.  530.)     The  small  generating-flask  (lOOc.c.)  is  fitted  with  a 
double  bored  rubber  stopper,  a  delivery  tube  6,  and  a  tube,  A,  which  connects 
with  a  hydrogen  generator.     In  the  generating-flask,  place  a  solution  of  20 
grams  of  potassium  hydroxide  in  40  c.c.  of  water  and  two  pieces  of  yellow 
phosphorus  as  large  as  a  bean  ;  now  pass  a  current  of  hydrogen  through  the 
flask  until  all  of  the  air  is  expelled,  shut  off  the  hydrogen,  and  then  heat  the 
generating-flask  in  a  sand-bath. t     Phosphine  will  pass  off,  and  will  take  fire 
spontaneously  when  it  reaches  the  air.     Demonstrate  this  by  placing  the  end  of 
your  delivery-tube  under  water  before  beginning  the  experiment;  the  individual 

*  Substances  frequently  must  be  heated  in  sealed  tubes.  The  end  which  is  closed  be- 
fore filling  the  tube  should  be  round,  like  the  bottom  of  a  test-tube,  and  as  thick  as  the  walls 
of  the  glass  tube.  After  the  substance  to  be  heated  is  added,  the  other  end  of  the  tube  is 
sealed  to  a  long  point,  not  by  heating  the  glass  and  drawing  out,  but  by  heating  and  allow- 
ing the  sides  to  fall  together,  so  that  the  sealed  point  is  as  thick  as  the  rest  of  the  tube. 
This  operation  requires  considerable  practice. 

t  A  shallow  iron  dish  containing  sand. 


530 


APPENDIX    OF 


bubbles  will  then  rise  to  the  surface  and  burn.  A  simpler  and  better  way  of 
preparing  phosphine  is  to  place  a  piece  of  calcium  phosphide  in  a  basin  of 
dilute  hydrochloric  acid ;  bubbles  of  phosphine  will  at  once  pass  off,  and  take 


Fig.  48. 


fire  spontaneously  when  they  reach  the  air.  The  gas  can  be  collected  in  a  test- 
tube  filled  with  water,  and  inverted  over  the  piece  of  calcium  phosphide. 
Phosphine  which  is  not  spontaneously  combustible  can  be  prepared  by  pass- 
ing the  gas,  generated  by  either  of  the  above  methods,  through  the  apparatus 
used  to  condense  sulphur  dioxide  (Note  43);  the  liquid  phosphine  which  is 
spontaneously  combustible  will  then  be  condensed,  while  the  gaseous  phos- 
phine will  pass  on. 

62.  THE  PREPARATION  OF  PHOSPHORUS  PENTOXIDE.  —  Place  a  piece 
of  phosphorus  on  a  porcelain  plate,  ignite  it,  and  cover  with  a  glass  bell ; 
phosphorus  pentoxide  will  collect  on  the  walls  of  the  latter  and  on  the  plate. 
Dissolve  some  of  the  pentoxide  in  water,  and  test  the  solution  with  blue  litmus ; 
expose  some  of  the  pentoxide  to  the  air;  to  the  solution  of  phosphoric  anhy- 
dride, add  silver  nitrate;  boil  another  portion  of  the  solution  for  some  time, 
nearly  neutralize  with  ammonia,  and  then  add  silver  nitrate.  Take  some 
secondary  sodium  phosphate  (ordinary  sodium  phosphate),  and  heat  a  little  of 
the  salt  in  a  hard  glass  test-tube  until  water  of  crystallization  is  driven  off; 
now  heat  more  strongly  (formation  of  pyrophosphate).  To  a  solution  of 
secondary  sodium  phosphate,  add  a  solution  of  calcium  chloride,  and  then 
perform  the  experiments  suggested  by  the  text  on  page  230.  Take  a  little 
sodium-ammonium  hydrogen  phosphate,  and  heat  in  a  hard  glass  test-tube 
(notice  the  odor!);  finally  heat  until  the  substance  forms  a  transparent  glass; 
place  some  of  this  transparent  glass  on  the  end  of  a  platinum  wire,  and  heat 
with  a  little  cobalt  nitrate,  manganese  chloride,  ferric  chloride ;  using  a  fresh 
drop  of  fused  sodium  metaphosphate  and  a  clean  wire  for  each  one  of  the  salts. 


LABORATORY   NOTES. 


531 


Phosphorus  trichloride,  when  dissolved  in  water,  produces  phosphorous 
acid.  Place  a  few  drops  of  the  phosphorus  trichloride  in  a  test-tube,  add 
water,  and  notice  how  the  oily  liquid  dissolves;  test  the  solution  by  means  of 
blue  litmus  paper;  evaporate  the  solution  to  dryness  on  a  water-bath  until  all 
of  the  hydrochloric  acid  has  passed  off,  and  then  add  an  alcoholic  solution  of 
iodine  (page  225,  paragraph  2,  c). 

Phosphorus  pentachloride,  when  dissolved  in  water,  produces  phosphoric 
acid ;  take  a  little  of  the  pentachloride  on  a  spatula,  and  add  it  to  water  in  a 
test-tube.  The  reaction  will  take  place  at  once,  with  a  hissing  noise;  the 
solution  of  phosphoric  acid  will  not  reduce  iodine  to  hydroiodic  acid,  as 
the  solution  of  phosphorous  acid  does. 

63.  MARSH'S  TEST  FOR  ARSENIC.  — The  accurate  details  of  this  method 
belong  in  works  especially  devoted  to  analytic  chemistry.  Take  your  hydro- 
gen generating-flask,  attach  a  U-shaped  drying-tube  to  this,  and  connect  the 
latter  with  a  hard  glass  tube,  which  has  been  drawn  to  a  point,  and  which  is 
constricted  at  two  places  by  being  drawn  out  in  the  flame  of  a  blast-lamp 
(Fig.  49). 

Fill  the  U-shaped  tube  with  granulated  calcium  chloride. 

Place  20  grams  of  pure  zinc  in  your  generating-flask,  add  dilute  sulphuric 
acid,  and  allow  a  brisk  current  of  hydrogen  to  traverse  the  apparatus ;  when 
all  is  safe,  ignite  the  jet  at  the  drawn-out  point  of  the  hard  tube,  place 
a  cold  porcelain  plate  in  the  flame,  and  see  if  the  flame  leaves  a  spot.  Now 
add  a  solution  of  arsenic  trioxide  in 
hydrochloric  acid  to  the  generating-flask 
through  the  thistle-tube;  in  a  few  min- 
utes arsine  will  be  developed,  and  the 
hydrogen  flame  will  assume  a  violet 
color,  with  a  white  smoke ;  now  heat  the 
hard  glass  tube  at  a  point  just  before  one 
of  the  constrictions;  amorphous  arsenic 
will  be  deposited  on  the  cold  portions  of 
the  tube  in  the  form  of  a  mirror;  this 
mirror  is  volatile,  and  can  be  driven  from 
place  to  place  along  the  tube  by  heating  it 
with  a  Bunsenburner.  Hold  a  cold  porce- 
lain plate  in  the  flame  at  the  tip  of  the 
hard  glass  tube;  this  will  cool  the  escap- 
ing gases  to  a  point  below  the  kindling 
temperature  of  arsenic;  the  latter  element  will  therefore  be  deposited  on  the 
plate  as  a  black  spot.  This  spot,  when  touched  with  a  drop  of  sodium  hypo- 
chlorite  solution  on  the  end  of  a  glass  rod,  will  be  instantly  dissolved 
(3  NaOCl  +  2  As  =  As.2O3  +  3  NaCl);  when  touched  with  nitric  acid,  it  is 
dissolved,  owing  to  oxidation  and  solution  of  the  arsenious  oxide  formed. 
Arsine,  when  passed  into  a  solution  of  silver  nitrate,  precipitates  metallic 
silver,  and  forms  arsenic  trioxide.  This  reaction  can  be  obtained  by  extin- 
guishing the  flame  at  the  tip  of  the  hard  glass  tube,  and  passing  the  mixture 


Fig.  49. 


532  APPENDIX   OF 

of  gases,  which  are  being  generated,  into  a  test-tube  containing  silver  nitrate 
solution ;  black,  metallic  silver  will  be  precipitated ;  filter  this  off,  and  add  am- 
monia to  the  clear  filtrate,  until  neutral;  yellow  arsenite  of  silver  will  be 
precipitated.  If  the  gas  which  is  passing  from  the  generator  is  stibine,  and 
not  arsine,  the  mirror  spots  on  the  tube  will  have  a  grayish  and  more  metallic 
appearance,  will  form  nearer  to  the  flame,  and  will  be  very  nearly  not-volatile. 
The  spot  on  the  porcelain  mirror  will  not  disappear  on  addition  of  sodium 
hypochlorite ;  it  will  turn  white  on  addition  of  nitric  acid,  because  the  oxide 
of  antimony  is  insoluble  in  that  substance.  If  stibine  is  passed  into  a  solution 
of  silver  nitrate,  black  silver  antimonide,  SbAg3,  is  precipitated;  if  this  pre- 
cipitate is  boiled  with  a  concentrated  solution  of  tartaric  acid,  then  the  anti- 
mony will  be  dissolved,  and  its  presence  can  subsequently  be  readily  proved  by 
precipitating  antimony  sulphide  by  means  of  sulphuretted  hydrogen.  For 
further  directions,  consult  some  work  on  analytical  chemistry.  (See  page  255.) 
(All  work  with  arsine  or  stibine  must  be  done  under  the  hood!) 

64.  A  QUICK  METHOD  FOR  DETECTING  ARSENIC.  —  Draw  a  piece  of 
hard  glass  tube  to  a  point,  as  is  shown  in  Fig.  50;  place  a  little  arsenious 
oxide,  or  a  small  particle  of  the  substance  which  you  suspect  to  be  arsenious 

oxide,  in  the  tip;  above  this, 
place  a  small  piece  of  char- 
coal ;  heat  the  tube  at  the  spot 
where  the  coal  has  lodged 
until  the  coal  is  red-hot,  and 
then  gradually  draw  the  tube 
through  the  flame  until  the 
arsenic  trioxide  is  heated ;  the 

f    50  latter  will  sublime,  and  will, 

in  passing  over  the  hot  char- 
coal, be  reduced  to  metallic  arsenic,  which  will  form  a  mirror  similar  to 
that  observed  in  Marsh's  test. 

The  experiments  which  can  be  performed  with  the  oxides  and  sulphides  of 
arsenic  will  suggest  themselves  during  the  study  of  chapters  xxxi.  and  xxxii. 

65.  STIBINE.  — The  preparation  of  stibine  is  exactly  like  that  of  arsine, 
so  that  the  pupil  may  follow  the  directions  given  in  Note  63,  excepting  that  he 
must  substitute  a  hydrochloric  acid  solution  of  antimony  trioxide  for  that  of 
arsenious  oxide.     The  distinctions  existing  between  the  spot  produced  on  a 
cold  plate  by  burning  arsine,  and  that  produced  by  stibine,  are  given  in  the  last 
part  of  Note  63;  the  experiments  showing  such  distinction  should  be  followed 
out.     As  antimony  and  arsenic  are  elements  which  are  commonly  met  with  in 
analytical  work,  and  as  their  reactions  are  fully  described  in  all  of  the  direc- 
tions for  qualitative  analysis  which  are  published,  it  is  scarcely  necessary  to 
enter  into  a  more  detailed  discussion  of  experiments  to  be  performed  with 
antimony  at  this  place.     By  following  the  text,  the  teacher  can  easily  select 
experiments  which  need  no  detailed  description  (pages  252  to  256,  and  chapter 
xxxiv. ).     The  same  is  true  as  regards  bismuth. 


LABORATORY   NOTES. 


533 


66.  DESTRUCTIVE  DISTILLATION  OF  COAL  AND  WOOD.  —  Unless  a  lab- 
oratory is  especially  well  appointed,  and  unless  the  pupil  has  plenty  of  time  at 
his  disposal,  it  will  not  be  possible  to  carry  out  the  dry  distillation  of  bitumi- 
nous coal  or  of  wood ;  when  such  work  is  attempted  it  belongs,  more  prop- 
erly, in  a  course  on  organic  chemistry.     The  pupil  should,  however,  place  a 
small  piece  of  wood  in  a  hard  glass  test-tube,  and  heat,  endeavoring,  as  much 
as  possible,  to  ascertain  the  nature  of  the  products  evolved ;  he  should  also  see 
if  the  gases  which  pass  off  will  burn.     The  same  experiments  should  be  per- 
formed with  a  piece  of  bituminous  coal. 

67.  THE  ABSORPTION  OF  COLORING  MATTER  BY  CHARCOAL.  —  Prepare 
a  solution  of  indigo  (Note  59),  place  in  a  300  c.c.  flask,  and  warm,  after  add- 
ing two  tablespoonfuls  of  animal  charcoal ;  *  filter,  and,  if  the  solution  is  not 
colorless,  repeat   the   operation.     Do  the  same  with  a  solution  of  iodine  in 
iodide  of  potassium.      The  absorption  of  gases  by  charcoal  was  described  in 
Note  49. 

69.  THE  FLAME.  —  The  fact  that  the  centre  of  a  flame  is  cold,  while  the 
outer  zone  is  hot,  can  be  demonstrated  by  turning  the  flame  of  a  Bunsen 
burner  down,  shutting  off  the  air  supply  so  that  the  flame  becomes  luminous, 
and  then  quickly  placing  a  sheet  of  filter  paper  upon  it  so  that  about  one-half 
of  the  flame  is  below  the  surf  ace.  of  the  paper  (Fig.  11,  page  283);  as  soon  as 
the  brown,  burnt  circle  appears  on  the  upper  side  of  the  paper,  withdraw  the 
latter  quickly. 

Diluting  illuminating  gas  with  a  non-combustible  gas  renders  the  flame  non- 
luminous.  —  Arrange  an  apparatus  as  is  shown  by  Fig.  51.  Attach  one  arm  of 
the  T-shaped  glass  tube  to  a 
tap  of  illuminating  gas,  and 
attach  the  other  to  an  appa- 
ratus generating  dry  carbon 
dioxide;  the  remaining  arm 
of  the  T-tube  is  fitted  with  a 
brass  pipe  which  can  be  heat- 
ed by  a  Bunsen  burner.  Turn 
off  the  supply  of  carbon 
dioxide  by  the  pinchcock, 
turn  on  the  gas,  and  light  the 
burner;  the  flame  will  of 

course    be    luminous.      Now  Fi 

turn   on   the   carbon  dioxide 

so  that  the  latter  gas  mingles  with  and  dilutes  the  illuminating  gas;  the  flame 
will  at  once  become  larger  and  non-luminous.  However,  if  the  brass  exit-tube 
which  serves  as  a  burner  is  heated  to  a  point  at  which  the  decomposition  of 
ethylene  into  methane  and  carbon  takes  place,  the  flame  will  once  more  be 
rendered  luminous,  although  still  diluted  by  carbon  dioxide. 


Animal  charcoal  can  be  obtained  at  any  chemical  supply  house. 


534 


APPENDIX    OF 


A  flame  can  be  extinguished  by  cooling  below  the  kindling  temperature.  — 
Light  a  Bunsen  burner,  and  then  place  a  piece  of  copper  wire  gauze  upon  the 
flame  so  that  the  latter   will 
about  be  bisected ;    the  flame    ji 
will  continue  to  burn  beneath 
the  gauze;  but  above  the  lat- 
ter, because  the  wire  conducts 
the  heat  away  too  rapidly,  no 
flame   will   be  seen.     The  re- 
verse of  this  experiment  (Fig. 
52)  is    shown  by  placing  the 
gauze  above  an  unlighted  bur- 
ner, turning  on  the  gas,  and 

then  lighting  above  the  gauze;  the  flame  will  then  not 
form  below. 

To  show  that  oxygen  will  burn  in  illuminating 
gas.  —  Construct  an  apparatus  as  shown  by  Fig.  53. 
The  tube  B  is  connected  with  an  illuminating  gas  tap, 
the  gas  turned  on,  and  allowed  to  run  until  all  air  is 
expelled  from  the  glass  bulb;  now  light  the  gas  at  Fig.  53. 

the  top,  the  double  bored  stopper  containing  tubes  A 

and  C  being  removed  for  the  purpose.  Attach  A  to  a  gasometer  which  will 
furnish  oxygen,  turn  on  a  slow  stream  of  that  gas,  and  then  bring  the  stopper 
which  holds  A  and  C  back  into  position.  As  the  current  of  oxygen  comes 
in  contact  with  the  burning  gas  escaping  from  the  bulb,  it  will  be  ignited,  and 
will  continue  to  burn  in  the  illuminating  gas. 

70.  THE  PREPARATION  OF  CARBON  MONOXIDE  BY  PASSING  STEAM 
OVER  RED-HOT  CHARCOAL.  —  Take  an  iron  gas-pipe  700  m.m.  in  length; 
attach  one  end  of  it,  by  means  of  a  rubber  stopper,  to  a  flask  of  300  c.c. 
capacity,  which  is  so  arranged  on  a  retort  stand  that  you  can  boil  water  in  it, 
and  that  the  steam  must  pass  through  the  iron  tube.  Connect  the  other  end 
of  this  iron  tube  with  a  safety  bottle  (Note  1),  and  put  a  delivery-tube,  bent  so 
that  it  will  open  under  the  water  in  your  pneumatic  trough,  into  the  bottle ; 
place  the  iron  tube  into  a  combustion  furnace  (Fig.  1,  Note  1),  having  previ- 
ously filled  the  tube  with  pieces  of  charcoal  which  are  broken  to  the  size  of  a 
pea;  heat  to  a  red  heat,  and  then  pass  steam  over  the  charcoal.  After  all  of 
the  air  has  been  expelled  from  the  apparatus  by  means  of  the  current  of  steam, 
collect  the  escaping  gas  over  the  pneumatic  trough  by  displacement  of  water. 
This  gas  will  be  carbon  monoxide  mixed  with  hydrogen,  as  can  be  proven  by 
inverting  one  of  the  cylinders  filled  with  it,  and  touching  a  lighted  taper  to 
the  mouth.* 


*  In  using  the  combustion  furnace  for  this  experiment,  take  care  that  the  iron  tube  is 
sufficiently  long  to  extend  some  little  distance  beyond  either  end  of  the  furnace;  otherwise 
the  rubber  stoppers  will  become  too  hot  and  will  melt.  You  should  feel  of  the  stoppers  from 
time  to  time;  and  if  there  is  danger  of  their  fusing,  cool  them  by  pouring  on  water. 


LABORATORY  NOTES. 


535 


71.  THE  PREPARATION  OF  CARBON  MONOXIDE  FROM  OXALIC  ACID.  — 
The  apparatus  is  the  same  as  that  used  for  nitrous  oxide  (Note  52,  Fig.  43). 
In  the  retort,  place  10  grams  of  crystallized  oxalic  acid,  add  60  grams  of  con- 
centrated sulphuric  acid,  and  heat  until  a  regular  evolution  of  gas  takes  place; 
pass  this  gas  through  a  wash-bottle  (Fig.  18)  containing  a  solution  of  potas- 
sium hydroxide  (one  part  potassium  hydroxide  to  two  parts  of  water)  before 
you  collect  over  the  pneumatic  trough.      By  means  of  this  wash-bottle  the 
carbon  dioxide,  which  is  generated  simultaneously  with  the  carbon  monoxide, 
is  absorbed  (it  forms  potassium  carbonate),  while  pure  carbon  monoxide  is 
collected;  the  latter  gas  will  burn  with  a  pale  blue  flame. 

72.  THE  PREPARATION  OF  CARBON  DIOXIDE.  —  The  apparatus  is  the 
same  as  that  used  for  the  preparation  of  hydrogen  or  of  hydrogen  sulphide 
(Fig.  34).     Charge  the  generating-flask  with  20  grams  of  marble  which  has 
teen  broken  to  the  size  of  a  hickory  nut,  pour  on  100  c.c.  of  hydrochloric  acid, 
-which  you  have  diluted  with  an  equal  amount  of  water,  and  pass  the  gas 
which  is  generated  through  a  wash-bottle  containing  water  (if  dry  carbon 
dioxide  is  required,  pass  the  gas  through  a  second  bottle  containing  sulphuric 
acid);  collect  the  carbon  dioxide  by  displacement  of  the  air  as  you  did  chlo- 
rine and  hydrobromic  acid. 

Take  a  number  of  other  carbonates  (sodium  carbonate,  potassium  carbon- 
ate, barium  carbonate,  etc.),  place  a  little  of  each  in  test-tubes,  and  add 
hydrochloric  acid  to  each  test-tube ;  ascertain  if  carbon  dioxide  is  given  off. 

73.  EXPERIMENTS  WITH  CARBON  DIOXIDE. — Fill  5  or  6  cylinders  with 
carbon  dioxide.  Construct  an  apparatus  such  as  is  shown  by  Fig.  54;  a  num- 
ber of  small  candles  are  placed  on  wires  which  are 
attached  in  an  upright  position  t$  a  larger  wire  which 
is  bent  in  the  form  of  a  flight  of  steps ;  light  all  of 
the  candles,  place  them  in  a  beaker  glass,  and  pass 
carbon  dioxide  into  the  latter  by  means  of  a  tube 
extending  to  the  bottom;  the  lights  will  be  extin- 
guished successively  from  below  upward ;  this  experi- 
ment will  also  demonstrate  the  fact  that  carbon 
dioxide  neither  burns,  nor  supports  combustion.  Pass 
some  carbon  dioxide  from  your  generator  into  a  solu- 
tion of  barium  hydroxide  and  into  one  of  calcium 
hydroxide :  filter  the  precipitates  formed  in  each  case, 
wash  them  from  the  filter  papers  into  test-tubes,  and 
add  hydrochloric  acid ;  prove  that  carbon  dioxide  is 
passing  off  by  holding  a  glass  rod  which  has  been  Flg  54 

dipped  into  lime-water  (a  solution  of  calcium  hydroxide)  just  within  the 
mouths  of  the  test-tubes;  if  carbon  dioxide  is  being  generated,  the  lime-water 
tvill  become  turbid.  Pass  carbon  dioxide  into  a  concentrated  solution  of 
potassium  hydroxide  and  into  one  of  sodium  hydroxide;  take  a  little  of  the 
solutions  so  formed,  and  add  hydrochloric  acid  to  them;  prove  that  carbon 
dioxide  passes  off.  Pass  carbon  dioxide  into  a  solution  of  lime-water  until 
the  precipitate  which  at  first  forms  is  re-dissolved  (formation  of  the  primary 


536  APPENDIX   OF   LABORATORY   NOTES. 

carbonate),  and  then  boil  the  solution;  add  hydrochloric  acid  to  the  solution 
before  it  is  boiled  and  to  the  precipitate  which  is  formed  by  boiling,  and  see 
if  carbon  dioxide  is  given  off  in  each  case. 

74.  THE  ETCHING  OF  GLASS  BY  MEANS  OF  HYDROFLUORIC  ACID.  —  Take 
a  shallow  lead  dish  3  inches  long,  2  inches  wide,  and  1  inch  deep;  put  about 
10  grams  of  powdered  calcium  fluoride  into  this,  and  cover  the  fluoride  with 
sulphuric  acid ;  prepare  a  piece  of  window  glass  by  dipping  it  into  melted 
paraffin,*  allowing  the  latter  to  cool,  and  etching  some  figures  by  scratch- 
ing away  the  paraffin  coating  with  the  point  of  a  knife ;  place  this  prepared 
glass  over  the  lead  dish  in  the  form  of  a  cover,  put  the  whole  into  a  warm 
place  under  the  hood,  and  allow  to  stand  for  6  or  8  hours;  the  glass  will 
then  be  etched  where  the  paraffin  was  scratched  away.  (In  working  with 
hydrofluoric  acid,  be  extremely  careful  not  to  allow  it  to  come  in  contact  with 
the  hands  nor  to  inhale  the  fumes.  The  ulcers  caused  by  hydrofluoric  acid 
burns  are  very  painful  and  slow  to  heal,  and  the  results  of  inhalation  of  the 
vapors  may  be  dangerous.) 

Preparation  of  silicon  tetrafluoride  and  of  Jluo silicic  acid.  —  Take  a  flask 
of  300  c.c.  capacity,  put  into  this  10  grams  of  quartz  sand  mixed  with  10 
grams  of  powdered  calcium  fluoride ;  fit  a  single  bored  rubber  stopper  to  the 
flask,  and  in  this  rubber  stopper  place  a  glass  delivery-tube  which  is  bent 
with  two  right  angles ;  the  end  of  this  delivery  tube  will  then  point  downward ; 
to  this  end  attach  a  funnel  by  means  of  the  stem  in  such  a  manner  that  the 
escaping  gas  must  pass  through  the  funnel,  and  place  a  beaker  of  water  under 
the  funnel  so  that  its  rim  just  touches  the  surface  of  the  water.  Pour  100  grams 
of  concentrated  sulphuric  acid  on  the  mixture  of  sand  and  calcium  fluoride  in 
the  generating-flask,  connect  the  apparatus  and  warm  gently;  silicon  tetra- 
flouride  will  pass  off,  and  will  come  in  contact'with  the  water  in  the  beaker;  by 
this  means  fluosilicic  acid  will  be  formed.  While  the  solution  of  the  latter  is 
being  formed,  it  will  at  the  same  time  become  filled  with  flakes  of  silicic  acid,  so 
that,  if  the  delivery-tube  had  not  been  widened  by  the  funnel,  it  would  soon  have 
become  clogged  and  an  accident  would  have  followed.  Filter  the  fluosilicic  acid 
from  the  precipitated  silicic  acid,  and  test  the  reaction  of  the  solution  towards  lit- 
mus; add  just  enough  potassium  hydroxide  solution  to  neutralize  the  fluosilicic 
acid  and  allow  to  stand;  nearly  insoluble  potassium  fluosilicate  will  separate. 

The  elements  which  follow  silicon  in  the  text  are  nearly  all  of  a  metallic  na- 
ture, and  none  of  the  experiments  which  will  fix  their  character  in  the  mind  of 
the  pupil  will  present  any  practical  difficulty.  It  is  better,  therefore,  for  the 
teacher  to  select  such  experiments  as  he  deems  proper  from  the  text;  indeed,  it 
often  seems  advisable  to  study  some  of  the  chemical  relations  of  the  metals  in 
the  laboratory  by  the  methods  of  qualitative  analysis,  while,  at  the  same  time, 
becoming  familiar  with  the  more  general  aspect  of  the  chemistry  of  those  ele- 
ments by  following  the  text-book  on  general  chemistry;  the  small  manuals 
containing  directions  for  qualitative  analysis  are,  however,  so  numerous  that 
it  seems  unnecessary  to  add  a  list  of  experiments,  which  practically  cover  the 
same  ground,  to  a  text-book  of  general  chemistry. 

*  Use  so-called  high  melting  paraffin. 


INDEX. 


The  numbers  refer  to  the  pages. 


Acetylene 


281        Acid,  phosphoric 95,  226,  232 


Acid,  Boric 330 

Bromic 130 

Carbamic 298,299 

Chlor-auric 409 

Chlorides  of  sulphur 157 

Chlorous 119 

Chlor-platinic 493 

Chlor-sulphonic 157 

Cyanic 297 

Cyanuric 297 

Disilicic 307 

Disulphuric 154 

Dithionic 155 

Ferric 483 

Fluoboric 330 

Fluosilicic 302,303 

Hydrobromic 80-82 

Hydrochloric 66-77 

Hydrocyanic 295 

Hydrofluoric 56,57 

Hydroiodic 85-87 

Hypobromous 129 

Hypochlorous 115,  119,  121 

Hyponitrous 208 

Hypophosphorous 231 


lodic 

Meta-arsenic  .  .  . 
Meta-arsenious  .  . 
Meta-antimonic .  . 
Meta-phosphoric 
Meta-silicic  .  .  . 
Meta-vanadic  .  .  . 

Nitric 

Nitrosyl-sulphuric   . 
Nitrous    . 


.  .  .  130,  132 
.  .  227,241,242 
.....  227 
255 

.  .  .  226,228 
227,  304,  305,  307 

441 

.     .     .       203-210 

147 

.     .    147,  148,  208 


Ortho-arsenic 227,  241,  242 

Ortho-arsenious 227 

Ortho-antimonic 255 

Ortho-phosphoric 227, 229 

Ortho-silicic 227,  304,  305 

Ortho-vanadic 441 

Pentathionic 155 

Perchloric 119,  126 

Per-iodic .130 


Phosphorous 95,  224 

Plumbic 323 

Pyro-antimonic 255 

Pyro-antimonous 254 

Pyro-arsenic 241 

Pyro-phosphoric 227,  231 

Pyro-vanadic 441 

Selenic 161 

Selenious 160 

Stannic  a 317 

Stannic  ft 317 

Stannous 317 

Sulpharsenic 245,246 

Sulpharsenious 244,  245 

Sulpho-dithio  carbonic 293 

Sulphuric,  56,  66,  80,  117,  136,  137,  145,  150 

Sulphurous 138,  158 

Telluric 161 

Tellurous 161 

Tetrathionic 155 

Thiosulphuric 154 

Trithionic 155 

Trisilicic 307 

Trithiocflrbonic   (see    sulpho-dithio 
carbonic). 

Acid  reactions 76,  177, 178 

Acids,  action  on  salts 142 

Affinity  for  bases 141 

Cause  of  character 177,  178 

Decomposition  by  metals    ....      31 

Definition  of   .     .    . 75 

Formation  of ,  from  anhydrides  .    .     116 

General  formulae  of 117 

Hydrated 131,  135,  145 

Nature  of 32 

Polybasic 139 

Relative  avidity  of 141,  142 

Strong  and  weak 141 

Uni-,  di-,  and  tribasic 139 

Valence  in 116,  117 

Affinity,  chemical     .......      10 

Agate 304 

Alabaster 418 

Alkali  metals,  carbonates  of  .    .    .    .389 


537 


538 


INDEX. 


Alkali  metals,  comparative  table  of   .    391 
Decomposition  of  water  by  .    .      383,  384 

Double  halides  of 388 

Halidesof 388 

Halides  of,  heats  of  formation  of     .    388 

Hydroxides  of 385,  386 

Oxides  o? 385 

Properties  of 384,  385 

Relations  of 383 

Sulphhydrates  of 386 

Sulphides  of 386 

Alkaline  earths,  carbonates  of   .     415,  416 

Chlorides  of 414* 

Comparison  of  properties  of  ...    412 

Metallurgy  of 411 

Oxides  and  hydroxides  of   .    .      412,  413 

Sulphates  of 416,  4}7 

Table  of 411 

Allotropism 48 

Alloys 249,250 

Aluminium,  basic  sulphates  of  ...    340 

Halides  of 336,  337 

Halides,  double  salts  of 337 

Hydroxides  of 30,  338 

Occurrence  of 333 

Phosphates  of      ........    340 

Preparation  of 333,334 

'Properties  of 334 

Trichloride 336 

Trichloride,  molecular  weight  of      .    336 

jTrioxide 338 

Trisulphide 341 

Alum 339,340 

Amalgams 191 

Amido  group 298 

Ammonia,  decompos'n  of,  by  chlorine,      63 

History  of 182 

In  atmosphere 170 

Preparation  of 183,  184 

Properties  of i.     .     .    .    185 

Solubility  of I    ...    187 

>    Table  of  compounds  of 263 

Volumetric  composition  of      .      185,  186 

Ammonia  liquor 184 

Ammonia-soda  process 390 

Ammonium  amalgam 191 

Carbamate 298 

Carbonate 298 

Cyanate 298 

Chloride 188,189 

Molybdate  .    .    . 452 

Nitrate 188 

Phospho-molybdate 454 

Sulphate 188,  189 

Chloride,  vapor  density  of   ....     189 
Ammonium  salts,  Decomposition  of 

184,  189,  190 
Formation  of 189,  190 


Ammonium  salts,  nature  of  ....    189 

Occurrence  of 182 

Table  of 263 

Anatas 438 

Anglesite 320, 417 

Anhydrides    .    14,  22,  25,  26,  30,  95,  130,  131 
Conversion  to  acids      .     115,  116,  117,  178 

Valence  in 115 

Anhydrite 417 

Antimonic  acids 255 

Antimonous  acids 254 

Antimony,  acids  of 254,  255 

Alloys  of ,..       249-251 

Basic  salts  of  . 253 

Double  salts  of 253 

Halides  of 252 

History  of 248 

Metallurgy  of 248 

Occurrence  of 248 

Oxides  of 254 

Pentachloride 181,  253 

Antimony,  pentasulphide  of    ....    256 

Pentoxide 254 

Properties  of 248,  249 

Sulphides  of 255,  256,  264 

Tetroxide    .' 254 

Tribromide 252 

Trichloride      .     .     .  181,  249,  252,  253,  263 

Trifluoride 252 

Tri-iodide 252 

Trioxide 249,254 

Trisulphide 255 

Apatite 211,418 

Arragonite 416 

Arsenates 242,  247 

Arsenic  acid,  oxidizing  action  of      242,  243 

Salts  of    . 241 

Table  of 247 

Arsenic,  halides  of 238,  239 

History  of 234 

Metallurgy  of 235 

Occurrence  of 234 

Properties  of 235 

Table  of  acids  of 247 

Table  of  oxides  of 247 

Table  of  sulphides  of  .     .    .     .      247,  264 

Pentafluoride 238 

Pentasulphide 245,246 

Pentoxide 239,  241 

Tribromide 238 

Trichloride 181,238,263 

Trifluoride 238 

Tri-iodide 238 

Trioxide 235,  239,  240,  241 

Trisulphide 244,  245,  247 

Arsenious  oxide 239, 247 

Poisonous  effects  of 240 

Chemical  action  of 241 


INDEX. 


539 


Arsenates 

Arsenites 241-217 

Arsenopyrite 2<*4 

Arsine,  preparation  of 236 

Properties  of 236,237 

Artificial  ice 187 

Atmosphere,  the,  ammonia  in 170 

Carbon  dioxide  in 167,  168 

Composition  of,  history  of  ....     165 

Pressure  of 170,  171 

Pressure  of,  measurement  of  .      171,  372 
Quantitative  composition  of    ...     166 

Relation  of,  to  life 174 

Solids  in 170 

Specific  gravity  of 173 

Water  vapor  in 169 

Atomic  heats  of  compounds,  rela- 
tions of      353,354 

Atomic  heats  of  elements,  relations 

of  352-354 

Atomic  hypothesis,  the     ...     4,  6,  347 
Atomic  volumes,  relations  of     .      365-367 

Atomic  weights,  absolute 7 

Determination  of 349-361 

Maximum 73,349 

Of  oxygen  family 106 

Standard  of 7 

Table  of 8 

Auric  chloride 408 

Hydroxide 408 

Oxide 408 

Sulphide 409 

Aurous-auric  oxide 406 

Chloride 408 

Aurous  chloride 408 

Cyanide 409 

Sulphide 409 

Avidity  of  acids     .    .  • .    .    .Ill,  142,  113 

Avogadro 70 

Avogadro's   hypothesis,    71,  72,   73,  348 

349,  360 
Azoimid 193,  194 

Barium,  isolation  of 412 

Properties  of 412 

Carbonate 292^415,416 

Chlorate 420 

Hydroxide  .  .  .  .  • 413 

Nitrate 420 

Oxide 413 

Permanganate 466 

Sulphate 416,417, 

Superoxide ill 

Barite 417 

Barometer,  height  of 171 

History  of 170 

Variations  of 173 

Barytocelestine 417 


14,25,26,30,31,177 

Neutralization  of 375-377 

Relative  strength  of 377 

Beauxite 333,334 

Benzine 279 

Beryll 307 

Beryllium>  isolation  of 411 

Properties  of 412 

Carbonate 415 

Chloride 414 

Hydroxide 413 

Oxide 413 

Sulphate 416,117 

Berzelius M 

Bismite 257 

Bismuth,  alloys  of 258 

Basic  halides  of 258 

Halidesof 259 

Metallurgy  of 257 

Native 257 

Occurrence 257 

Oxides  of 259 

Properties  of 258 

Salts  of 260,  261 

Hydroxide 259,  260 

Monosulphide 261 

Monoxide 259 

Nitrate 260 

Nitrate,  basic 260 

Pentoxide 259 

Subnitrate  (see  basic  nitrate). 

Sulphate 261 

Telluride 257 

Tetroxide 259 

Tribromide 258 

Trichloride 258 

Trifluoride 258 

Tri-iodide 258 

Trioxide 257,  259 

Trisulphide 257,261 

Bismuthinite 257 

Blast  furnace,  construction  of   ...    473 

Changes  in 474 

Borax  . 328,  331 

Berates,  ortho-,  meta-,  and  tetra-    .    .    331 

Occurrence  of 328 

Boric  acid,  occurrence  of 330 

Ortho-  and  meta 331 

Properties  of 331 

Borocalcite 328 

Boron,  acids  of 329 

Occurrence  of 328 

Preparation  of 328 

Properties  of 329 

Hydride 329 

Nitride 332 

Oxy-chloride 332 

Phosphate 331 


540 


INDEX. 


Boron,  trichloride 329 

Trifluoride 329 

Trioxide 330 

Boron  family,  elements  of 326 

Oxides  of  elements  of 326 

Valence  of  elements  of  .  .  >  .  .  326 

Halidesof '  .  .  327 

Braunite 460 

Bromides,  formation  of 82 

Bromine,  hydrate  of 80 

Occurrence  of 79 

Oxy-acids  of 129,  130 

Preparation  of 79,  80 

Properties  of 79 

Monochloride 132 

Water 80 

Brookite 438 

Butane  (see  diethyl). 

Bunsen  burner,  the 283 

Cadmium,  alloys  of 428 

Metallurgy  of 427 

Occurrence  of 426 

Properties  of 424 

Carbonate 429 

Chloride 429 

Hydroxide 428 

Oxide 428 

Sulphate 429 

Sulphide 430 

Calcite 292,416 

Group  of  minerals 416 

Calcium,  isolation  of 412 

Properties  of 412 

Carbonate 292,  413,  415,  416 

Carbonate,  primary 418    | 

Chlorate 125,420 

Chloride 414,420 

Dimanganite 465 

Hydroxide 30,  44,  413 

Hypochlorite 122,  123 

Oxide 413 

Phosphate 212,418 

Phosphate,  primary  ....  229,418 
Phosphate,  secondary  .  .  .  230, 418 

Phosphate,  tertiary 229 

Silicates 420,  421 

Tungsten 445 

Calcium  and  magnesium  carbon- 
ate     292,416 

Calomel  (see  mercurious  chloride). 

Carbamic  acid 298,  299 

Carbon,  allotropic  forms  of     .    .       269-273 

Amorphous 272,273 

Halidesof 285,286 

Hydrogen  compounds  of  .  .  274-284 
Hydrogen  compounds  of,  table  of  .  279 
.Nitrogen  compounds  of  ...  294-299 


Carbon,    nitrogen    and   oxygen    com- 
pounds of 297-299 

Nitrogen  and   oxygen    compounds 

of,  table  of  . 299 

Occurrence  of <;69 

Oxides  of 286-291 

Table  of  compounds  of 319 

Carbonates 291 

Natural,  table  of 292 

Primary  and  secondary   .     .     .-     291,  292 

Carbon  dioxide  . 22 

Changes  in  atmosphere    .    .     .      167,  168 

History 289 

In  atmosphere 166-168 

Preparation 290 

Properties 290,291 

Carbon  disulphide 95, 292 

Carbonic  acid,  meta- 291,  299 

Ortho 285, 286 

Carbonic  acids 267 

Carbon  monoxide 286 

Preparation 286, 287 

Properties 287 

Poisonous  action 288 

Tetrabromide 286 

Tetrachloride 75,  285 

Carbon  family,  acids  of 267 

Elements  of,  comparison  of  ...  265 
Hydrogen  compounds  of  .  .  265,  266 
Hydrogen  compounds  of,  table  of  .  266 

Oxides  of 267 

Oxides  and  acids  of,  table  of   ...    268 

Sulphides  of,  table  of 268 

Table  of  compounds  of 325 

Carbonyl  chloride     .    .    .    .288,  289,  298 

Cassiterite 312,438 

Cavendish 27,  36, 165 

Celestine 417 

Cement 414 

Cerium,  compounds  of 439 

Halidesof 440 

Occurrence  of 439 

Oxides  of 440 

Cerussite 320 

Chalcedony 304 

Chalcocite 397 

Chalcopyrite 91,  397 

Charcoal 273 

Chemism 10 

Chlorates,  formation  of 124 

Properties  of 125 

Chlo-rauric  acid 409 

Chloric  acid 119 

Decomposition  of 126 

Properties  of 126 

Chlorides,  conversion  of,  into  oxides      112 

Formation  of 76,  77 

Formulae  of 63,  64 


INDEX. 


541 


Chlorides,  valence  of  elements  in  .    .    109 

Of  sulphur 156 

Chlorine,   action   on   hydrogen    com- 
pounds   63 

Bleaching  action 64 

Combustion  in 63 

History  of 58 

Occurrence 58 

Oxides  of 119,120 

Oxidizing  action  of 65 

Oxy-acidsof 119,128 

Preparation  of 59,  60,  61,  62 

Properties  of 61,  62 

Chlorine  dioxide 119,  127 

Hydrate 62 

Monoxide     .......   115,  119, -120 

Trioxide 119,  127 

Water 64 

Chloroform  (see  methiu  chloride). 

Chlorous  acid 119 

Chlor-platin  amines 494 

Chlorsulphonic  acid 157 

Chromates 449,  450,  451 

Chrome-alum 447 

Chromic  acid 448 

Chloride 447 

Hydroxide 446 

Oxide 446 

Chromite 444 

Chromites 446 

Chromium,  Acid  chlorides  of      .      448,  449 
Compounds  of,  preparation      .      451,  452 

Compounds  of,  uses    , 452 

Occurrence  of 444 

Preparation  of 445 

Properties  of 445,  446 

Chromium  family,  elements  of      .    .    443 
Elements  of,  comparison     ....    443 

Halides  of 444 

Tables  of  compounds  of  .     .   443,  458,  488 

Chromous  chloride 451 

Hydroxide 451 

Chromous-chromic  oxide     ....    451 

Chromspinell 444 

Chromyl  chloride 448 

Cinnabar 427, 435 

Clay 340 

Purification  of 341 

Coal,  formation  of 168,  271,  272 

Cobalt,  metallurgy  of 477 

Occurrence  of 472 

Properties  of 472 

Sulphides  of 472 

Cobalt  amines    .    .    .    .  ^ 486 

Cobaltglance 234 

Cobaltic  nitrate 485 

Oxide 485 

Cobaltite 472 


Cobaltous  chloride 484 

Nitrate 485 

Oxide 484 

Sulphate 417 

Sulphide 485 

Columbates 441 

Columbite 441 

Columbium,  compounds  of     ....    441 

Occurrence  of 441 

Combustion 22,  23,  24 

In  chlorine 24 

In  oxygen 22,  23 

Rapid  and  slow 23 

Constancy  of  matter,  law  of     .    .    .  2, 15 

Copper,  alloys  of 400 

Chlorides  of 401,402 

Metallurgy  of 398 

Occurrence  of 397 

Oxides  of 401,  402 

Properties  of 398 

Salts  of 401,402 

Sulphides  of 404 

Copper,  silver,  and  gold,  alloys  of  .    400 

Properties  of 397,  400 

Resemblance  to  alkalies       ....    396 
Table  of  compounds  of     ...      409,  410 
Corrosive    sublimate   (see   mercuric 
chloride). 

Corundum 333 

Crocoite 444 

Cryolite 333,  334 

Cupric  carbonates,  basic ....      403,  404 

Chloride 402 

Hydroxide 260,  402 

Nitrate 403 

Oxide 38,402 

Sulphate 42, 402, 417 

Sulphate,  formation  of    ...      137,  402 

Sulphide 404 

Cuprous  chloride 401 

Iodide 401 

Oxide 401 

Sulphide 404 

Cyanates 297,298 

Cyanic  acid 297,298 

Cyanamide 299 

Cyanides 294 

Cyanogen 294, 295 

Dalton,  John 3, 69 

Dalton's  hypothesis 3 

Davy 37,  58,  104 

Deliquescence 43 

Dialysis 305 

Diamond 269, 270 

Diaspor 333 

Dichromates 450 

Diethyl 278,279 


542 


INDEX. 


Dimethyl 278,279 

Disilicates :  .  .  307 

Dissociation 61,  351 

In  solution 351,  380-382 

Disulphuric  acid 154 

Dithionic  acid 155 

Dolomite 292,416 

Double  decomposition 57 

Laws  of 379,380 

Double  salts,  structure  of 337 

Dulong  and  Petit,  law  of  ...  351-356 
Dumas 38,  165 

Efflorescence 43 

Ekaboron 373 

Elements,  division  of 12 

Electro-negative 13,368 

Electro-positive 13,  368 

Number  of 9 

Polyvalent,  compounds  of   .     .     .    .     114 

Prediction  of 373 

Relation  between 17 

Specific  heat  of,  relation  of     ...    352 
Table  of  molecular  weights  of     .    .    426 

Electrolysis 13, 37 

Energy,  chemical    .    10,  11,  23,  24,  46,  74,  76 

Kinetic 11,  12 

Potential     . 10,  11 

Transformation  of 11 

Ethane  (see  Dimethyl). 

Ethylene 280 

Decomposition  of 282,  283 

Properties  of 281,282 

Structure  of 280 

Equations,  chemical   ....  16,  19,  20,  74 

Equivalent  quantities 141 

Weights 39,  357,  358 

Euxenite     .    .    .    .   • 437 

Faraday's  law 359 

Ferric  acid 483 

Chloride 482 

Hydroxide 30,  481 

Oxide 481 

Sulphate 482 

Sulphide 482 

Ferrites 481,483 

Ferrous  carbonate 292,480 

Chloride 479 

Columbate 441 

Hydroxide 29,  260,  479 

Oxide 479 

Selenide 103 

Sulphate 42,417,479,480 

Sulphide 97,480 

Tantalate 441 

Tungstate 445 

Ferrous-ferric  oxide     ...     22, 438,  483 


Flame,  the 22,  282,  283 

Flint 304 

Fluorine,  combustion  in 55 

History  of 53 

Occurrence  of 55 

Preparation  of 55 

Properties  of ,      55 

Valence  of 303 

Fluosilicic  acid,  constitution  of     .    .    303 

Preparation  of S02 

Properties  of  .......      303, 304 

Formic  aldehyde 167 

Formulae,  chemical 15,  IP,  20 

Structural 108,  111,  112,  113 

Gadolinite  .    .    . 437,439 

Earths 437,  442 

Galenite      91,320 

Gallium 343,  344 

Garnet 307, 333 

Gases,  calculation  of,  to  standard  vol- 
ume   172,173 

Kinetic  theory  of 68,  69 

Relation  of  pressure  and  volume  of,    172 
Ratio  between  specific  gravity  and 

combining  weight  of      ....      68 

Relation  by  volume 70,  71 

Relation  of  specific  gravities  to  mo- 
lecular weights  of       ....     72,  73 

Gay  Lussac 37,  68,  69,  165 

Law  of 68 

Germanium 309,  310,  311 

Chloroform 310 

Dioxide 310 

Disulphide 311 

Monosulphide 311 

Monoxide 310 

Tetrachloride 310 

Gersdorfite 473 

Glass 420,421,422 

Glucinum  (see  beryllium). 

Gold,  alloys  of 400 

Chlorides  of 408 

Hydroxides  of 408 

Metallurgy  of 399,  400 

Occurrence  of 398 

Oxides  of 408 

Graphite,  occurrence 270 

Uses 271 

Greenockite 426 

Guanidine 299 

Guano 418 

Gunpowder 391 

Gypsum 417 

Halhydric  acids,  relative  stability  of 

53,  54,  88 
Heats  of  formation  of      ....  87,  105 


INDEX. 


543 


Halhydric  acids,  tables  of  .    .    87,  105,  266 

Halogens,  the 53 

Anhydrides  and  acids  of      ....     134 

Comparison  of 53,  54 

Compounds  with  each  other  .  132,  133 
Oxy-acids  of,  comparison  .  .  179,  384 
Oxy-acids  of,  heats  of  formation  of,  162 
Oxy-acids  of,  nomenclature  of  .  .  119 
Oxy-acids  of,  tables  of  ...  119, 133 

Hausmannite 460 

Hematite     . 472 

Heptane 279 

Hexane 279 

Hornblende 307,333 

Humboldt 37,  165 

Hydrargyllite 333 

Hydrated  acids 131, 135 

Nomenclature  of 227 

Hydration 131 

Hydrazin 192,  193 

Hydro  bromic  acid 80,81 

Heat  of  formation  of 87 

Hydrocarbons  in  coal-oil      ....    279 
Hydrochloric  acid     ......    fifi,  fiz. 

Action  on  metals 76 

Action  on  oxides 76 

Action  on  hydroxides 76 

Heat  of  formation  of  .  l .  .  .  .  74,  87 
Volumetric  composition  of  .  .  .  68,  72 

Hydrocyanic  acid  - 295,  296 

Hydrofluoric  acid 56,  57 

Structure  of 303 

Hydrogen 27 

Combustion  of 35 

Diffusion  of 34 

History  of 27 

Nascent 137,  207 

Occlusion  of 34 

Occurrence  of 27 

Preparation  of      ....       28,  29,  31,  35- 

Properties  of 33,  35 

Hydrogen  and  oxygen,  ratio  between 

the  atomic  weights  of    ...       8,  39 
Hydrogen  antimonide     ....      251, 263 

Arsenide 236,  263 

Boride 329 

Hydrogen  compounds  of  not-met- 

als,  comparison  of    .  90,  176,  177,  178 
Comparative  tables  of      ...      175,  176 

Decompositions  of 98 

Heats  of  formation  of      ....  98,  108 

Structure  of 108 

Unsaturated 109 

Valence  in 108,  109 

Hydrogen  hy peroxide 50,  51 

Oxidizfng  action  of 51 

Hydrogen  persulphide 101 

Phosphides 214,  217,  263 


Hydrogen,  selenide 103 

Silicide 301 

Hydrogen  sulphide 96,  100 

Composition  of 99 

Decomposition  of 98, 99 

Heat  of  formation  of 98 

Preparation  of '     96,  97,  183 

Reactions  of 100 

Hydrogen  Telluride 104 

Hydroiodic  acid 85,  86 

Heat  of  formation  of 87 

Hydroxides 29,30,43 

Decomposition  of,  by  heat  ....    260 

Formation  of 260 

Of    metals  and  not-metals,  resem- 
blances between 157 

Hydroxyl    .     .    30,  31,  43,  117,  128,  131,  140 

Hydroxylamin 191 

Hydrochloride 192 

Hyperoxides 25 

Nature  of 323 

Hypobromous  acid 129 

Hypochlorites,  formation  of  ....    121 

Decomposition  of 122,  123 

Changes  of,  by  heat 124 

Hypochlorous  acid 115,  121 

Decomposition  of 123 

Oxidizing  action  of     ....       123,  124 

Properties  of 123 

Salts  of 121 

Hypophosphorous  acid 232 

Illuminating  gas,  formation  of  .      183,  282 

Indium,  properties  of 344 

Compounds  of 345 

lodates 132 

lodic  acid 130, 132 

Iodine,  history  of 83 

Oxy-acids  of 130 

Preparation,  properties  of  ...     83,  84 

Iodine  monobromide 133 

Monochloride 133 

Pentafluoride 133 

Pentoxide 131 

Trichloride 133 

Iodine  —  starch 50 

Iridium 489,  490,  492 

Iron 31 

Alloys  of 478 

Cast,  gray 474 

Cast,  white 474 

Chemically  pure 476 

Metallurgy  of 473,  474 

Occurrence  of 472 

Passive 477 

Properties  of 476 

Spiegel 475 

Sulphides  of 31 


544 


INDEX. 


Iron,  wrought 475 

Iron,  cobalt,  and  nickel,  comparison 

of 470 

Properties  of 470 

Relation  of,  to  periodic  system      470,  471 

Table  of 488 

Iron  pyrites 91,472,484 

Isomorphism 42,  356 

Application  to  atomic  weights      356,  357 

Examples  of 357,  358 

Law  of .    356 

Kaolin 318,  333,  340 

Lampblack 273 

Lanthanum 437 

Lavoisier 2, 15,  27,  36, 165 

Law  of  definite  proportions  .  .  3,  347,  348 
3Iultiple  proportions  ....  3,  5,  347 

Lazurite 397 

Lead,  metallurgy  of 320 

Occurrence  of 320 

Oxides  of 322 

Properties  of 321 

Red  oxide  of 324 

Salts  of 322 

Table  of  compounds  of 325 

Acetate 323 

Carbonate 323 

Chloride 322 

Chromate 322,  444,  450 

Dioxide 323 

Hydroxide 322 

Monoxide 322 

Molybdate 444 

Nitrate 201 

Sulphate 322, 417 

Sulphide 91,324 

Tungstate 445 

Le  Blanc  soda  process  ....     389,  390 

Leucite 307 

Light,  chemical  action  of     ...      406,  407 

Ligroine 279 

Lime  (see  calcium  oxide). 

Limonite 472 

Magnesite 416 

Magnesium,  isolation  of 412 

Properties  of 412 

Magnesium-ammonium  arsenate  419,  420 
Phosphate 419 

Magnesium  carbonate .415 

Chloride 61,  415 

Hydroxide 29,  76,  413 

Oxide 413 

Phosphate 419 

Pyrophosphate 419 

Secondary  phosphate 419 


Magnesium,  sulphate     ....      416, 417 

Magnetite 472,483 

Malachite 397 

Manganates 465 

Manganese,  alloys  of 460 

Compounds  of 460, 469 

Occurrence  of 459, 4GO 

Oxides  of 460 

Properties  of 460 

Relation  of  compounds  of    ....    468 

Relation  of,  to  periodic  system    .     .    459 

Tables  of  compounds  of  ...      468,  469 

Manganese  dioxide  19,  20,  25, 60, 79, 84, 460 

Chemistry  of 460,  463 

Hydroxides  of 465 

Occurrence  of 460,  463 

Manganite 460,  462 

Manganites 465 

Manganic  alum 463 

Chloride 462 

Hydroxides 462 

Oxide 460,  462 

Sulphate 462 

Manganous  carbonate 462 

Chloride 461 

Hydroxide 461 

Oxide 460 

Sulphate 461 

Sulphide 462 

Manganous-manganic  oxide  460,  461,  4.63 

Markasite 91,  472 

Marsh  gas  (see  methane). 

Mass  action,  laws  of 378,  379 

Melaconite 397 

Mendelejeff 361,373 

"Men  dele  Jeff's  table  of  the  periodic 

system 371 

Mercuric  chloramide 434 

Chloride 432,  433 

Cyanide 294,435 

Iodide 434 

Nitrate 435 

Oxide 19,  431 

Sulphide 435 

Mercurous  chloramide    ....      432,  434 

Chloride 431 

Iodide 434 

Nitrate    .     . 434 

Oxide 430. 

Sulphide 435 

Mercury,  metallurgy  of  ......    427 

Occurrence  of 427 

Oxides  of 430 

Properties  of 425 

Salts  of 431,435 

Metals,  characteristics  of    ....     12,  13 

Oxides  of 14 

Metaphosphoric  acid    ....     228,233 


INDEX. 


545 


Methane,  action  of  chlorine  on    ...  276 

Occurrence  of      .     .     * 274 

Preparation  and  properties  of      .     .  275 

Table  of 279 

Volumetric  composition  of  ....  276 

Methin 277 

Chloride 277 

Methyl 277 

Methyl  chloride 277 

Methylen 277 

Methylen  chloride 277 

Meyer,  Lothar 361,  367 

Mica 307 

Millerite 473 

Minium 324 

Mitscherlich 356 

Mixtures 5 

Moissan 55 

Molybdates 453,454 

Molybdenites 444,452 

Molybdenum,  acids  of 453 

Halidesof 453 

Hydroxides  of 453 

Occurrence  of 444 

Oxides  of 452 

Preparation  and  properties  of     .    .  446 

Sulphides  of 454 

Molybdenum  dioxide 452 

Disulphide 444 

Monoxide 452 

Pentachloride 453 

Molybdic  acids 453 

Molybdite 444 

Molecular    weights,    determination 

of 349 

Molecular  weights  of  gases  .    .    .    72,  73 

Molecules 9 

Endothermic  and  exothermic      .     .  12 

Stability  of 11 

Muscovite 333 

Nascent  action 51,  137,  207 

Neutralization,  laws  of  ....  375-382 
Relation  to  dissociation  .  .  .  381,  382 

Neutralization  of  acids,  heat  of  .  143,  376 

375 
Laws  of 375,  377 

Nickel,  alloys  of 478 

Cyanides  of 487 

Metallurgy  of 477 

Occurrence  of 473 

Properties  of 478 

Sulphate 417 

Sulphides  of 473 

Nickelous  chloride 487 

Cyanide 487 

Hydroxide 486 

Niobium  (see  columbium). 


Nitrates,  decomposition  of 208 

Formation  of 204,208 

Nitric  acid,  anhydride  of 202 

History  of 203 

Preparation  of 204 

Production  of 202,  203 

Properties  of 205,206 

Reduction  of  ....  183,  198,  199,  207 
Saltsof 208 

Nitric  bxide,  oxidation  of  .    .   198,  200,  201 

Preparation  of 198,  199 

-    Properties  of 200 

Nitrites 208 

Nitro  group 147 

Nitrogen,  acids  of 195 

History  of 163,  164 

In  atmosphere 166 

Occurrence  of 164 

Oxides  of 180,  195 

Preparation  of 164 

Properties  of 165 

Nitrogen  dioxide 198,201 

Formation  of 202 

Properties  of 202 

Nitrogen  oxides,  heats  of  formation  of    210 
Table  of 209 

Nitrogen  pentoxide 202 

Trichloride 180 

Trioxide 147,  148,  201 

Nitrogen  family,  comparison  of  mem- 
bers   176 

Halides  of,  comparison  of    .     .      180,  181 

Halides  of,  table  of 263 

Hydrogen  compounds  of  176,  177,  263,  266 
Oxides  of,  comparison  of  .  .  180,  263 
Sulphides  and  sulpho-salts,  table  of, 

263,  264 

Table  of  atomic  weights  of .     .      175,  262 

Table  of  hydrogen  compounds  of     .     175 

218,  263 

Table  of  oxides  and  acids  of,  179,  263,  264 
Table  of  properties  of  ...  175,  262 
Table  of  specific  gravities  of  vapors 

of 262 

Valence  in  oxides  of 115 

Nitrosyl-sulphuric  acid,  decomposi- 
tion of  148 

Formation  of 147 

Nitrous  oxide 196 

Composition  of 197 

Effects  of 197,  198 

Preparation  of 196 

Properties  of 196,  197 

Not-metals,  characteristics  of     .    .     12,  13 
Oxides  of 14 

Occlusion  of  gases 34 

Oligoclase 308 


546 


INDEX. 


Olivin 307 

Orpiment 234 

Orthoclase 308 

Ortho-phosphoric  acid      .    .    .     229,233 

Salts  of 229 

Osmium 489 

Dioxide  of 492 

Oxides  of 490 

Osmosis 305 

Osmotic  pressure 306 

Osteolite 211,418 

Oxide  of  manganese,  black  (see  man- 
ganese dioxide). 

Oxides 21,24 

Oxides,  classification  of 25,  26 

Formulae  of,  and  valence  of     .      112,  113 

Nomenclature  of 26 

Valence  in 110 

Oxygen 18 

Combustion  in 22 

Discovery  of 18 

In  atmosphere 166 

Occurrence  of 18 

Preparation  of 18,  19,  20 

Properties  of 21 

Oxygen  family,  atomic  weights  of     .    106 

Comparative  table  of 105 

Comparison  of 89, 90 

Hydrogen  compounds  of,    89,  90,  105,  266 

Melting  points  of 105 

Oxides  and  oxy-acids  of 135 

Specific  gravities  of 105 

Specific  gravities  of  gases    ....     105 

Valence  in  oxides  of 115 

Oxy-hydrogen  blow-pipe     ....     36 

Ozone 47 

Composition  of 48 

History  of 47 

Oxidizing  action  of      ....    50,  51,  52 

Preparation  of 47,  49,  56 

Properties  of 49 

Palladium 33,  489,  490 

Cyanide 489 

Dioxide N   .     .    492 

Palladium-hydrogen 33,207 

Pentathionic  acid 155 

Pentane 279 

Pentlandite 473 

Pentoxide  of  iodine     .....      130,  131 

Nitrogen 203 

Phosphorus 22 

Perchlorate  of  potassium, 

Decomposition  of 125 

Formation  of 125 

Properties  of 126 

Perchloric  acid 119 

Properties  of 120 


Per-iodates 132 

Per-iodic  acids 130 

Hydrated 131 

Periodic  system ,  the  ....  361-374 

Arrangement  of 362,  363 

Relations  of  atomic  volumes  to  365,  366 

Relations  of  families  in 364 

Relations  of  halides  to  .....  370 

Relations  of  hydrides  to 370 

Relations  of  oxides  to  ...  370,  372 
Relations  of  valence  to 369 

Permanganic  acid 466 

Anhydride 460,  466 

Petroleum 279 

Ether 279 

Phlogiston 18,27 

Phosgen  (see  carbonyl  chloride). 

Phosphates,  ortho,  meta,  and  pyro  .  228 

229,  231 

Phosphine,  liquid 217 

Preparation  of 214 

Properties  of 215,  216 

Volumetric  composition  of  ....  215 

Phosphites 227 

Phosphonium  compounds  .  .  .  216,  215 
Iodide 217 

Phosphorescence 213 

Phosphoric  acids,  95,  214,  220,  221,  224,  226 

Nomenclature  of 222, 226 

Oxidation  of 225 

Reactions  of 226 

Salts  of 225 

Phosphorus,  acids  of,  table  of  ...  233 

Allotropic  forms  of .  212 

Halides  of,  table  of 219 

Halides  of",  reactions  of 220 

History  of 211,212 

Hydrogen  compounds  of  ....  214 

Occurrence  of 211 

Oxidation  of 214 

Oxides  of,  table  of  ....  223,  231,  233 

Oxy-halides  of 221 

Poisonous  action  of 213 

Red 213 

Yellow 213,214 

Phosphorus  oxybromide 221 

Oxychloride 221 

Pentabromide 219,  220 

Pentachloride  .  .  .  181~  219,  220,  263 

Pentafluoride 219,  220 

Pentasulphide 95 

Pentoxide 22,214,223 

Tribromide 81,  219,  220 

Trichloride  .  .  75,  81,  181,  219,  220,  263 

Trifluoride 219 

Tri-iodide 85,219 

Trioxide 223 

Trisulphide 95 


INDEX. 


547 


Photography 406 

Pitchblende A.    .    .    445 

Plaster  of  Paris 418 

Platinic  chloride 492,493 

Platinous  chloride 495 

Platinum,  cyanides  of 489 

Dioxide  of 492 

Oxides  of 490 

Platinum  group,  comparison  of      .    .    489 

Cyanides  of 489 

Halidesof 491 

Occurrence  of 491 

Properties  of 490,  491 

Utensils  of 495 

Plumbago  (see  graphite).. 

Polysulphides 155, 386 

Structure  of 387 

Porcelain 341 

Potassium 30 

Bromate 129 

Brom-aurate 409 

Chlorate 20,  120 

Chlorate,  formation  of 124 

Chlorate,  decomposition  of      ...     125 

Chlor-aurate 409 

Chlorite 120 

Chromate 449,465 

Bichromate 450 

Ferricyanide 297 

Ferrocyanide 296 

Fluosilicate 304 

Hydroxide 30,  43,  387 

Hypobromite 129 

Hypochlorite    ....     120,  121,  122,  123 
Hypochlorite,  decomposition  of,  by 

heat 124 

Hyponitrite 208 

Iodide 50 

lodo-aurate 409 

Manganate 465 

Nitrate 391 

Pentamanganite 465 

Permanganate 467, 468 

Permanganate,  decomposition  of     .    467 
Permanganate,  oxidizing  action  of,    467 

Perchlorate 120,  125,  126 

Sulphides 386,387 

Uranate. 457 

Precipitate,  red 18,  19 

Priestley 18,  27,  182,  196,  201 

Primary  salts,  change  to  secondary,  140, 152 

Prince  Rupert's  drops 422 

Pyrolusite 460 

Pyr  ©phosphoric  acid     .    .    .    .     231, 233 
Pyrosulphuric  acid  (see  disulphuric acid). 

Quartz 304 

Raoult .350 


Raoult's  Law \     350, 351 

Reactions,  endothermic  and  exother- 
mic      12 

Realgar 234 

Recrystallization 41 

Reinite 445 

Rhodium 489 

Cyanides  of 489 

Dioxide  of 492 

Oxides  of 490 

Rutile 438 

Saltpetre,  Chile 204,  390 

Salts,  formation  of   ...      25,  30,  32,  77,  78 

Acid  and  neutral 140 

Basic 252 

Primary  and  secondary HO 

Reactions  of 141 

Scandium,  chemistry  of 437 

Predicted  properties  of    ...      373,  437 

Scheele 58 

Scheelite 445 

Secondary  salts,  change  to  primary,  140, 152 

Selenious  acid 160 

Selenites,  primary  and  secondary    .    .    160 
Selenium,  allotropic  forms      ....     102 

Dioxide  of 160 

History  of 102 

Isolation  of 102 

Monosulphide  of 161 

Occurrence  of  ... t 102 

Properties  of 103 

Trioxideof 161 

Siderite    .  ' 292,  472 

Silicates,  di 307 

Disintegration  of 340 

Meta 267,307 

Ortho 267,  307 

Tri 308 

Silicic  acids,  di- 307 

Dialyzed 305 

Meta- 304 

Ortho- 304 

Tri- 307 

Silico-chloroform 301 

Silicon,  occurrence  of 300 

Preparation  of 300 

Properties  of 301 

Halides  of 301,  302 

Hydride 301 

Oxides  of 304 

Dioxide 304 

Dioxide,  amorphous 304 

Tetrachloride :XX),,301 

Tetrafluoride 301,  302 

Silver,  alloys  of 400 

Halides  of 405 

Halides  of,  uses  in  photography  .     .    406 
Metallurgy  of 398,  399 


548 


INDEX. 


Silver,  occurrence  of    .  ....    398 

Oxides  of 405 

Chloride 405 

Hyponitrite 209 

Nitrate 407 

Sulphide 407 

Slag 473 

Smaltite 234,472 

Sodium 29 

Acetate 275 

Sodium  carbonate     43,  290,  291,  292,  389,  390 

Preparation  of 389,  390 

Chloride .    ....  7ft,  142 

Dimolybdate 453 

Hydroxide 29 

Hyposulphite  (see  sodium  thiosulphate). 

Metaphosphate 229 

Metavanadate 441 

Nitrate 204,390 

Octomolybdate 453 

Orthovanadate 441 

Phosphate,  primary 229 

Phosphate,  secondary 229 

Pyrovanadate 441 

Silicate 304 

Sulphate 142 

Sulphate,    action    of    hydrochloric 

acid  on 142 

Sulphate,  primary 152 

Sulphate,  secondary 152 

Tetramolybdate 453 

Thiosulphate 154,406 

Trimolybdate 453 

Tungstate    ;• 455 

ions 41 

;cific  gravities  of  gases     ...    72,  73 

jpe,  the 393 

jpic  analysis 393-395 

im ,  the 393 

Absorption 395 

Sphserocobaltite 472 

Spinells,  the     333,  339,  438,  444,  447,  460,  481 


Stannic  acid,  ortho 
Meta    . 


267 

267 

a 317 

0 317 

a  and  ft,  table  of 319 

Stannic  chloride 315,  317 

Double  salts  with 316 

Oxide 316 

Sulphide 318 

Stannous  acid 315 

.Chloride 313,  314,  433 

'Hydroxide 315 

Oxide 316 

Sulphide 318 

Status  nascendi 51,  137 

Steel    .  .     .    . 476 


Steel,  preparation  of 475,  476 

Stibine    49 251 

Stibnite 248 

Stibionyl 253 

Stoichiometric  quantities     ....       6 

Stoltzite 445 

Strass 421 

Strontianite 292,416 

Properties  of 412 

Isolation  of 412 

Strontium  carbonate 292,  416 

Chlorate 420 

Chloride 415 

Hydroxide 413 

Nitrate 420 

Oxide 413 

Sulphate 416, 417 

Structure  of  hydrogen  compounds,    108 

Substitution 29,277 

Sulphantimonites 255 

Sulphantimonates 256 

Sulpharsenates 245,  246,  247 

Sulpharsenites 244,  245,  247 

Sulphates,  primary  and  secondary,   152,  153 

Acid  and  neutral 152 

Sulphides,  formulae  of 94,  <J5 

Formation  of 100 

Solubility  of 100 

Sulphites,  change  by  heat 139 

Sulphocarbonates  (see  thiocarbonates). 

Sulphoferrites 483 

Sulphostannates 318 

Sulphur,  allotropic  forms  of    ....      93 

Crystalline  forms  of 93 

Halidesof   ., 156 

History  of 91 

Isolation  of 92 

Occurrence  of 91 

Oxides  and  oxy-acids  of  ...      134,  156 
Oxy-  and  sulpho-  acids  of,  table  .     .     159 

Plastic 94 

Properties  of 93 

Rare  oxides  of     ...     ./.'   ...     158 
Resemblances  to  chlorine    ....      95 

Sulphur  dioxide 22,  117, 134 

History  of 136 

Occurrence  of       135 

Oxidation  of 138 

Preparation  of 136,  137,  148 

Properties  of 137,  138 

Sulphur  trioxide 117, 134,  144 

Preparation  of 144 

Properties  of 144 

Dissociation  of 145 

Sulphur    family,    comparison     with   • 

/  chromium  family 443 

Heats  of  formation  of  compounds 

of.  162 


INDEX. 


549 


Sulphur   family,    oxides    and   acids 

of 135,  162 

Sulphuric  acid 117 

Action  of,  on  metals 137 

Chambers 149 

Commercial  preparation  of .     .      148,  149 

Concentration  of 150 

Constitution  of 145 

Continuous  process  of  manufacture 

of 147 

Heat  of  solution  of      ....      151,  152 

History  of 146 

Hydrated 145,  151 

Hydrated,  salts  of 153 

Hydration  of 150,  151 

Manufacture  of 147 

Properties  of 150 

Reduction  of 136,  151 

Uses  of,  in  laboratory 152 

Sulph'hydrates,  formulae  of    ....     94 
Sulphuretted  hydrogen    ....   96-101 

Sulphurous  acid 139 

Change  of,  on  heating 139 

Oxidation  of    ,m 139 

Structure  of 158 

Sulphuryl-chloride 138,  157 

Decomposition  of 115 

Sun,  chromosphere  of 395 

Superphosphates 418 

.Jiymbols,  chemicul 14tJLfi,  19,  20 

Tantalates 441 

Tantalum 441 

Tellurates 161 

Telluric  acid 161 

Tellurium 104 

Occurrence  of 104 

Properties  of 104 

Dioxide 160 

Disulphide 162 

Trisulphide 162 

Tetradymite 257 

Tetrahedrite 397 

Tetrathionic  acid 155 

Thallium 345 

Chlorides  of 346 

Oxides  of 345 

Properties  of 345 

Sulphides  of 346 

Theory,  dualistic 13 

Thiocarbonates 293 

Thionyl-chloride 157 

Thionic  acids,  constitution  of     ...  156 

Thiosulphates 154,  155 

Thorium 439 

Tin ]    '  312 

Dioxide 316 

Halidesof 314,315 


Tin,  history  of 312 

Hydroxides  of 315 

Metallurgy  of 312 

Occurrence  of 312 

Oxides  of 316 

Properties  of 313 

Sulphides  of 318 

Table  of  compounds  of 319 

Monoxide 316 

Tin-stone  (see  cassiterite). 

Titanium 438,439 

Torricelli 171 

Tridymite 304 

Trisilicates 308 

Trithionic  acid 155 

Troilite 472 

Tungsten 445 

Acids  of 455 

Halides  of 454,  455 

Occurrence  of 445 

Preparation  of 445 

Properties  of 446 

Dioxide 454 

Hexachloride 455 

Trioxide 444,  445 

Tungstic  acids 455 

Tungstite 445 

Unsaturated  hydrogen  compounds,  108 

Uranates 457 

Uranic  oxide 456 

acid 457 

Uranium,  occurrence  of 445 

Oxides  of 456 

Preparation  of 

Properties  of 

Uranous  oxide 4 

Uranous-uranic  oxide  ......  4 

Uranyl 456 

Chemical  relations  of 457 

Salts  of 456 

Hydroxide 456,  457 

Urea 298,299 

Valence 107 

Of  elements  towards  chlorine  .     .     .    109 
Of  elements  towards  hydrogen    .  107,  108 

281 

Of  elements  towards  oxygen    ...     Ill 
Of    elements    towards    polyvalent 

ones no,  112 

Law  of 111,112,113,114 

Relation  to  atomic  weights      .      114,  369 

Varying  , 115 

Vanadates 441 

Vanadinite 440 

Vanadium,  acids  of 441 

Chlorides  of ,         .441 


550 


INDEX. 


Vanadium,  comparison  with  phospho- 

rus and  arsenic      ......  441 

Occurrence  of      ........  440 

Oxides  of    ..........  440 

Vitriols,  the    ..........  417 

Vivianite     ...........  211 

Volumes  of  gases,  law  of  combina- 

tion of  ..........  G& 

"Water,  chemical  history  of      ...     30,  37 
Composition  of,  by  volume      .    37,  72M3 
Composition  of,  by  weight  ....   f  38 

Crystallization  of    ......     42,  43 

Decomposition  of,  by  chlorine     .    .      63 
Decomposition  of,  by  metals   .    29,  30,  31 
Energy  of  formation  of    .....      45 

Formation  of  ........     34,  36 

Properties  of  .........      40 

Quantitative  composition  of    ...      39 
Solution  in  ..........      41 

Vapor  in  atmosphere  ......     169 

"Waters.,  natural  .........     45 

Pollution  of     .........      45 

Mineral   ............      45 

Wavellite  ...........    340 

Witherite  .........     292,416 

Wollastonite 

"Work 

Measure  of 

Wulf  enite 

Wurtzite    ....... 


307 

10 

11 

444 

426 


Ytterbium 
Yttrium  . 


437 
437 


Zinc 

Alloys  of     .     . 

Hydroxide  of  . 

Metallurgy  of  . 

Occurrence  of. 

Oxide  .... 

Properties  of  . 
Zinc  blende     . 


32 

......  427 

.  .  ' .  .30,  76,  428 

427 

426 

428 

428 

...     91,  344,  426 
Zinc,  Cadmium,  Mercury,  compari- 
son of    423,  424 

Compounds  of 423 

Relations  of,  to  periodic  system  .    .    423 
Specific  gravities  of  vapors  ....    426 

Table  of  compounds  of 436 

Zincates 428 

Zinc  carbonate 429 

Chloride      . 76, 429 

Sulphate 417, 429 

Sulphide 91,  97,  430 

Telluride 104 

Zincite 426 

Zircon 438 

Zirconium,  compounds  of 439 

Occurrence  of      ...    .     .^     .     .    438 
Tetrachloride v*?T    .    .     .    439 

Appendix  of  Laboratory  Notes,  497  to  536 


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